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Q NBS SPECIAL PUBLICATION 557

U.S. DEPARTMENT OF COMMERCE /National Bureau of Standards

Chemical Kinetic Data Needs

for Modeling

the Lower Troposphere
NATIONAL BUREAU OF STANDARDS

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Chemical Kinetic Data Needs for Modeling

the Lower

, Troposphere + sbtcioJL: Dubr, _i

I

Proceedings of a Workshop

held at Reston, Virginia

May 15-17, 1978

John T. Herron, Robert E. Huie,

and Jimmie A. Hodgeson, Editors

National Measurement Laboratory

National Bureau of Standards

Washington, DC 20234

U.S. DEPARTMENT OF COMMERCE, Juanita M. Kreps, Secretary

Luther H. Hodges, Jr., Under Secretary

Jordan J. Baruch, Assistant Secretary for Science and Technology

NATIONAL BUREAU OF STANDARDS, Ernest Ambler, Director

Issued August 1979

Sponsored in part by

Environmental Protection Agency

Research Triangle Park, NC 2771
Library of Congress Catalog Card Number: 79-6001 25

National Bureau of Standards Special Publication 557

Nat. Bur. Stand. (U.S.). Spec. Publ. 557. 105 pages (Sept. 1979) COOEN: XNBSAV

U.S. GOVERNMENT PRINTING OFFICE

WASHINGTON: 1979

For sale by the Superintendent of Documents, U.S. Government Printing Office, Washington, D.C. 20402

Stock No. 003-003-02111-3 Price $4

(Add 25 percent additional for other than U.S. mailing)
Foreword

It is increasingly recognized that the integrity of the data input

is one of the most serious limiting factors in modeling complex chemical

systems. The reliability of the results generated in modeling studies

is critical considering their application to environmental regulation

and control

.

The purpose of this workshop was to bring together modelers, chemical kineticists, theoreticians and program managers, in order to define the critical data needs for modeling the troposphere. This collection of review papers, comments, and recommendations should serve a wide community

of atmospheric scientists in identifying and attacking priority problem

areas.

The National Bureau of Standards is pleased to be responsible for this publication, and to have joined with the Environmental Protection Agency as cosponsors of the workshop.

John D. Hoffman, Director

National Measurement Laboratory

National Bureau of Standards
Preface

It used to be said, "At least the air we breathe is free." For about

30 years that has not been the case in the United States. We are now paying about 10 billion dollars a year in an increasingly difficult

effort to maintain our atmosphere at a degree of pollution near current

levels. That sum will almost certainly increase as we turn towards alternate energy sources such as coal, shale oil and crude oils with

high sulfur contents. It will grow even larger as we learn more about

the impacts of various chemical substances, both natural and man made,

on the complex ecology in which man exists on this planet.

By contrast one can estimate that about 3 million dollars will be spent this year (1973) on basic research, by all government agencies for learning the molecular details of the chemistry of air pollution. It is perhaps very flattering to the scientific community involved in this effort to consider that such a sum will suffice to make significant

progress in the understanding of the chemistry of air pollution. But the scientists involved share no such illusions. In 30 years of research effort we have learned a great deal about the chemistry of our "dirty"

atmosphere but we can hardly pretend to give quantitative answers to questions which are increasingly being asked, such as, "What will the effect be on our atmosphere of removing or adding X tons per week of substance A?" Yet important economic decisions rest on the answers to such questions. The sums spent on basic research are dwarfed by the sums now being spent on regulation and abatement. Research funds are

the least expensive aspect of the problem of air pollution, and without

the answers that they might provide, our expensive efforts may be largely

squandered.

The present conference brings together many of the leaders in the basic research effort directed towards air pollution. They have examined

the scientific details with a fine microscope and come up with what must

seem to the layman an endless multitude of unanswered questions. Some

of these may never be answered and some of them need urgently to be resolved. Air pollution will not go away. It will become worse even with current efforts. No city on earth will escape its effects. If we hope to ameliorate it, we must devote a more profound and longer range effort to its understanding than we have done so far. The recommendations

of this symposium point the direction these efforts must take. As chairman, I would like to take this opportunity to express my appreciation

to my fellow colleagues who have given generously of their time and effort towards making this a fruitful meeting.

Sidney W. Benson University of Southern California

Los Angeles, CA 90007
Abstract

This is a report of the proceedings of a workshop on chemical kinetic

data needs for modeling the lower troposphere, held at Reston, Virginia,

May 15-17, 1978. The meeting, sponsored by the Environmental Protection Agency and the National Bureau of Standards, focussed on six key problem

areas in tropospheric chemistry: reactions of olefins with hydroxyl radicals and ozone, reactions of aldehydes, free radical reactions, reactions of oxides of nitrogen, reactions of aromatic compounds, and reactions of oxides of sulfur.

The report includes a summary and list of major recommendations for further work, review papers, discussion summaries, contributed comments, recommendations, and an attendance list.

Key words: Aldehydes, aromatics, chemical kinetics, data needs, free radicals, modeling, N0

x

, olefins, S0

x

, troposphere.

In order to describe experiments adequately, it has been necessary to identify commercial materials and equipment in this book. In no case does

such identification imply recommendation or endorsement by the National Bureau of Standards, nor does it imply that the material or equipment is necessarily the best available for the purpose.

v

Contents

Page

Foreword iii Preface iv Abstract v Introduction ix

Summary and Recommendations of Workshop 1

Session I: Reactions of Olefins with Ozone and Hydroxyl Radicals

An Evaluation of Chemical Kinetic Data Needs for Modeling

the Lower Troposphere: Reactions of Olefins with Hydroxyl

Radical and with Ozone — Hiromi Niki 7 Summary of Session 14 Comments 15 Recommendations 23

Session II: Aldehydes

Tropospheric Chemistry of Aldehydes — Alan C. Lloyd 27

Summary of Session 46 Comments 46 Recommendations 47

Session III: Organic Free Radicals

Organic Free Radicals -- David M. Golden 51

Summary of Session 61 Comments 62 Recommendations 66

Session IV: N0

X Chemistry

Tropospheric Chemistry of Nitrogen Oxides - A Summary of the Status of Chemical Kinetic Data -- Richard A. Cox 71 Summary of Session 74 Comments 74 Recommendations 79

Session V: Aromatics

Reactions of Aromatic Compounds in the Atmosphere -- Dale G. Hendry 85

Summary of Session 91 Comments 92 Recommendations 95

vii
Contents (continued)

Session VI: S0

x Chemistry

S0

x Chemistry (Abstract only) -- Jack G. Calvert 99

Summary of Session 99 Comments 100 Recommendations 101

Workshop Attendees 103 Subject Index 107 Author Index 107

viii
Introduction

The Environmental Protection Agency and the National Bureau of Standards' Office of Environmental Measurements and Center for Thermodynamics and Molecular Science sponsored a workshop entitled "Chemical Kinetic Data Needs for Modeling the Lower Troposphere," at Reston, Virginia, May 15-17, 1978. The objective of the workshop was to assess and make recommendations on mechanistic and kinetic data needs for modeling chemical transformations occurring in the lower troposphere.

The workshop was organized around six major topics: reactions of olefins

with hydroxy! radicals and ozone, the chemistry of aldehydes, free radical chemistry, the chemistry of oxides of mitrogen, the chemistry of aromatic

compounds, and the chemistry of the oxides of sulfur. These general topics cover almost all of the important problem areas in homogeneous chemical kinetics of interest to the atmospheric scientist. Heterogeneous processes

were not included.

Each technical session opened with a review paper followed by a discussion

period. A set of recommendations was prepared based on the review paper and subsequent discussions.

This report of the meeting includes the review papers (with one exception), discussion summaries, written contributions to the discussion, and recommendations. It opens with an overview and summary of the workshop recommendations.

We want to thank all those involved in organizing and running the workshop,

and all those who through their participation helped make it a success.

ix

Summary and Recommendations of Workshop

The recommendations of the workshop are given

in detail following each session. Here we

summarize the major themes of the workshop and list

the most important areas requiring additional

experimental or theoretical work.

Major advances in the chemistry of the troposphere depend on understanding the chemistry of

large molecules and free radicals. Real atmospheres contain large amounts of hydrocarbons Ci, and

greater, aromatic compounds, natural products such

as terpenes, and large aldehydes, ketones, phenols,

etc, which are their photooxidation products, as

well as a large class of oxygen containing free

radicals which are the intermediates in these

photooxidation reactions. These complex molecules

are not only involved in the NO-NO conversion

process, but almost certainly are important precursors to atmospheric aerosols. We will not understand either oxidant or aerosol formation until we

attack the problem of large molecules.

Clearly, moving away from fairly simple

surrogate reactants such as propylene, and

considering the whole range of atmospheric

pollutants creates a problem in scale. There are

far too many molecules and potential reactions to measure everything. A proper attack on this

problem involves a judicious mix of experiment and

theory. We may illustrate this by considering

one of the most pressing problems - the chemistry

of alkoxy radicals. Large alkoxy radicals can

isomerize, decompose, or react with oxygen (as well

as NO and SO , see below). The relative rates of

thesex process^s must be known. Even though this

problem was extensively considered at this workshop it is doubtful if a convincing solution was

given. Measurements are needed to provide a base

set to allow for the development of theoretical

estimation schemes.

Another class of reactions of great importance

both in the atmosphere and in laboratory investigations are alkylperoxy radical reactions. In the

absence of NO, these radicals react with themselves

to produce aldehydes and alcohols. There is considerable uncertainty as to the mechanism of

these reactions. The suggestion made at the

meeting that one of their products might be the

Criegee intermediate emphasizes the need for much

more work in this area. At the same time the

peroxy radicals formed in the reactions of OH with

olefins in the presence of 0 2 need identifying.

In addition we need data on acetyl and acetyl - peroxy type radicals. There are questions as to formation of acids, particularly from formyl

radical reactions, which cannot be explained on

the basis of existing data.

Another major deficiency is in the area of

reactions of aromatic compounds. Recognition of their importance is fairly recent. In particular

we need to know about rates and mechanisms of reaction of aromatics with OH radicals. This

will involve extension of existing experimental

approaches and development of new ones. The branching ratios for different products need to

be measured and the subsequent chemistry of these products needs to be considered.

The possibility of making the Criegee inter- mediate from alkylperoxy radicals was noted above. The Criegee intermediate is presumably a primary

product of an ozone-olefin reaction. Its subsequent fate is of great importance. A crucial question is whether it decomposes or is stabilized,

and if stabilized what chemical reactions it can undergo. There is some evidence that small Criegee intermediates decompose. For the large ones however, there is very little quantitative data. This question needs resolution since the Criegee intermediate has been postulated to be a potential oxidizer for NO, S02 , olefins, etc.

In addition to treating free radical reactions

in terms of i somen zati on, scission, self-reaction,

and reaction with 0 2 , we must consider reactions with NO and SO . Here we are faced with problems

of rates' and mechanisms and in particular the problem of the role of association reactions. For peroxy radicals the starting point is the H02 -N0 reaction. The new value for the rate constant has

had a dramatic effect on the models. It needs to be studied over a wide range of conditions

(temperature, pressure) to confirm this value under atmospheric conditions. The use of the rate constant data for H0 2 + NO for RO2 + NO reactions

may be invalid. Direct measurements are needed.

In addition, for large alkylperoxy radicals we need

to know if alky! nitrates are products since this

is a chain terminating reaction.

Similar considerations apply in the case of the reactions of alkylperoxy radicals with N0 2 , although it is not likely that the peroxynitrates

formed in simple association reactions would have

a significant lifetime in the atmosphere. How- ever, if the reaction can lead to an aldehyde and nitric acid it could be of considerable importance.

Reactions of alkoxy radicals with NO and N02 can also proceed via channels leading to adducts

or to HNO or HONO respectively. The overall rate constants and branching ratios need to be determined. Similar considerations apply to OH reactions with NO and N0 2

.

A somewhat different approach to the question

of the importance of adduct formation is to consider the thermal stability of peroxy radicals. Work of this kind has been done for some PAN

1

Session I

Reactions of Olefins with Ozone and Hydroxyl Radicals

AN EVALUATION OF CHEMICAL KINETIC DATA NEEDS

FOR MODELING THE LOWER TROPOSPHERE: REACTIONS OF OLEFINS WITH HYDROXYL RADICAL AND WITH OZONE

Hiromi Niki

Ford Motor Company

Dearborn, Michigan 48121

Needs for improved kinetic and mechanistic data for the reactions of olefinic hydrocarbons

with hydroxyl radical and with ozone have been evaluated from the view point of modeling the chemistry of the lower troposphere. Research priorities for removal of various uncertainties

in these reactions have been discussed briefly.

Keywords: Hydroxyl; kinetics; olefin; ozone; review; troposphere.

1. Introduction

In planning for the abatement and control of air

pollution, it is essential to establish the quantitative chemical relationship between source

emission and the resulting air quality. At present,

concerted modeling efforts are being made to

achieve this goal. Chemical interpretation of smog

chamber data [I] 1

, prediction of "ozone-isopleth"

[2], and air-shed modeling [3] of Los Angeles

Reactive Pollutant Program (LARPP) data [4] are a few of the notable examples of such endeavor.

Clearly, the degree of success of modeling work is governed, in large part, by the availability of

reliable kinetic data. Despite recent progress,

the knowledge of the chemical reactions taking

place in the lower troposphere is far from

satisfactory. There exist numerous critical

uncertainties in the kinetics and mechanisms of

reactions involving a large variety of atmospheric

constituents. This paper is intended to assess

the needs for improved experimental data for

olefin reactions with HO and with 0 3 .

For the purpose of evaluating the existing needs

for improved kinetic and mechanistic information

for these reactions, their potential role in the

lower troposphere is discussed briefly. Further,

the current knowledge of these reactions is illustrated by some of the work published within

the last few years. This paper is not intended to

be an extensive literature review, but rather, to convey the author's thoughts on the future

direction and research priorities in the area of tropospheric chemistry. In view of the urgency

and long-term interest in establishing firm

scientific bases for the abatement of air pollution

problems, the more systematic data evaluation

efforts, e.g., the "chemical reaction matrix"

method [5] used for the evaluation of the stratospheric chemistry, should be made in the future.

2. Atmospheric Role of Olefin Reactions

with HO and with 0 3

.

Olefins are among the most reactive classes of organic compounds present in the lower troposphere.

In particular, their possible role in the formation

of photochemical smog has been well-recognized

over the last three decades [6]. As a result, olefins have been used extensively as surrogate

hydrocarbons in laboratory smog studies, and have played the crucial role in the development of smog chemistry [1]. It now appears that the atmospheric

fate of olefins are governed primarily by their reactions with HO and with 0 3 . Conversely, these reactions are responsible for regulating the atmospheric concentrations of HO and 0 3

. The latter aspect is of primary importance, since HO is the major chain carrier of atmospheric reactions

and determines the role of other hydrocarbons and organic compounds in the formation of "oxidant", e.g., 0 3

. Whether 0 3 -olefin reactions can lead to the formation of "excess" 0 3 or alternatively serve

as a sink for 0 3 is another key question to be answered.

It must be stressed that a quantitative evaluation of the atmospheric role of the olefin reactions, or for that matter any other reactions,

can be made only on the basis of numerical modeling

studies for a given source distribution and strength under a variety of meteorological

conditions. Clearly, relative importance of HO and

0 3 reactions involving various olefins can vary markedly between "fresh" and "aged" air masses because of the chemically and meterol ogi cal ly induced changes in relative and absolute olefin concentrations

.

Figures in brackets indicate literature references

at the end of this paper.

To illustrate the relative importance of various olefinic and other types of hydrocarbons, table 1 shows the results obtained by Calvert [7] for the relative rates of HO-radical attack on hydrocarbons

7
Table 1. Estimated relative rates of HO-radical attack on hydrocarbons and CO present

0823 sample (LARPP Operation #33, Smog #2, 425-ft (AGL) Flight). 3

Compound

(RH), ppb,

mol basis

Rel . rate,

HO reaction Compound

(RH), ppb,

mol basis

CH,, 2.010 2.8

C 2 H 6 49 2.2

*C 2 H^ 43 11.6

C 2 H 2 38 1.0

C 3 H 8 37 4.4

*C 3 H 6 8.7 21.8 *iso-C l4 H 10 12 4.2 n-C^Hio 37 13.3 l-C^Hg 1.5 8.9

iso-C^He 3.0 27.9

iso-C 5 H 12 44.3 29.2

n-C 5 H 12 16.2 8.9 Cyclo-C 5 H 10 2.6 2.1 *1-C 5 H 10 4 ^ 2.4 *2-Methylbutene 0.8 ^ 7.4 *2,2-Dimethylbutene 0.8 -v 4.7

2- Methyl pentane 11.0 9.1

3- Methylpentane 10.0 8.3

*1 -Hexene 1.7 'v 10.0 n-Hexane 10.0 7.1 *Cyclohexene 10.7 ^ 10.7 2,2,3-Trimethyl butane 7.7 4.4

C 6 H 6 8.2 6.3

6.9 6.3

3-Methyl hexane

*1 -Heptene 0 n-C 7 H 16 Methyl cycl ohexane

2.2.3- and 2,3,3- Trimethyl pentane

2. 2. 4- Tri methyl pentane

Tol uene

*1 -Methyl eye lohexene

2.2.5- Trimethyl hexane

n-C 8 Hi 8 EtC 6 H s p.m-Xylenes

o-Xyl ene

n-C 4 H 2(j n-PrC G H 5 sec-BuC 6 H 5 n-C i nHoo

2-Methyl hexane

n

" c i 2 H 26 COTota l alkane

Total _ al kene

Total aromatic hydrocarbons

Grand total

6.3

4.4

4.3

3.7

2.5 2.0

4.7 1.0

2.1 4.1

1 .4 6.0

1.3

1 .0 5.0

1.1

1 .0 0.3

1 ,910

in the

Rel . rate,

HO reaction

5.8 ^ 25.9

3.8

4.7

1 .4

1 .8 17.6

<v 6.0 0.9

2.2 ^-4.2 ^32.1

11.3

1 .6 ^ 1.2 W.O

1.5

1.5 0.5 47.8 (12.1%)

128.0 (32.5°Q 138.3 (35.1%) 79.7 (20. 2

CQ

393.8

fa Taken from table IV in ref. [7]. Asterisks indicate olefins which contribute significantly to the removal of HO-radical.

in Los Angeles air samples [4]. It should be noted in this table, that a variety of olefins, indicated by asterisks, contribute significantly

to the removal of HO-radical. The olefins as a whole are responsible for 35 percent of the HO removal by hydrocarbons. Similarly, relative

removal rates of hydrocarbons by 0 3 can be estimated for the data given in table 1. The results are shown in table 2. Only the olefinic hydrocarbons are listed in this table, since 0 3 reactions with both paraffinic and aromatic hydrocarbons are negligibly slow. Notably, in this particular case, by far the dominant hydro- carbon-ozone reactions involve cyclic olefins,

i.e., cyclohexene and 1 -methyl cycl ohexene rather

than the straight chain olefins commonly used as surrogates. Calvert [7] also estimated the comparable fractional rates of removal of C 3 H 6 by

0 3 (3.2 percent h" 1

) and by HO (4.8 percent h" 1

) for conditions given in table 1. These examples

amply demonstrate that detailed hydrocarbon

analyses of air samples are essential to the understanding of atmospheric chemistry. Because of the close chemical coupling among 0 3 , NO and N0 2 concentrations via photo-stationary

relationship [9], the role of olefins in the oxidant formation can be assessed only in terms of their effects on NO chemistry. Specifically, the key question is "how do the atmospheric reactions

of olefins with HO and 0 3 control NO and N0 2 con- centrations?" The relevant chemistry is shown schematically in figure 1, where R represents H atom or hydrocarbon radicals formed, for instance, from olefin reactions with HO and 0 3 . Thus, to

Table 2. Estimated relative rates of 0 3 attack

on the olefins given in table 1.

(RH) Rel . 0 3 a Rel . removal'

' Compound p_p_b rate rate (%)

C 2 H„ 43 .15 2.3

C 3 H 6 8.7 . 1.0 3.2 1-C„H B 1.5 1.6.,,! .5 iso-C„H, 3 1.0 1.2

1- C 5 H 10 4 .8 1.2 2-methylbutene .8 37.9 11.2 2,2-dimethylbutene .8

b 1.1 .3 1-hexene 1.7 .9 .5 cyclo-hexene 10.7 13.0 51.4 1-heptene 4.4

b .9 1.4

1 -methyl cycl ohexene 4.7

b 15.4 26.7

|*Taken from ref. [8]. Estimated from analogous reactions.

8
Fig. 1. Reactions controlling 0 3 level

in the lower troposphere.

account for the behavior of 0 3 , it is necessary to know the kinetics and mechanism of various RO and

I R0 2 species interacting with NO and all other

I labile atmospheric constituents? In fact, the (predictive capability of modeling work on 0 3

I formation hinges on the understanding of free

radical chemistry involving NO [1]. It is fair to state that at present, there is a serious lack of reliable information on the question "how effectively do olefins generate radical species

and what is the subsequent fate of these radicals

in the atmosphere?" In the following sections

j the current knowledge on olefin reactions with HO

land 0 3 will be discussed to address these questions,

j 3. Current Knowledge of Olefin Reactions

With HO and With 0 3 .

A. HO-olefin reaction

Kinetics : As noted in the preceeding section,

the atmospheric life time and radical formation efficiency of olefins are governed, in large part,

by the HO reactions. Therefore, these rate constants should be determined with utmost

accuracy. In recent years, numerous direct and relative measurements of these rate constants have

been made over a wide range of temperatures and diluent pressures [10]. In particular, various

direct experimental methods employing flash photolysis-resonance absorption and fluorescence, and discharge-flow-laser magnetic resonance have been successfully used to measure the decay rates of

HO in the absence of interferences from secondary

(reactions. Thus, the overall accuracy of the

rate constants determined by these methods is determined by the inherent signal-to-noise ratios

in the HO decay curves and by the measurement of

the reactant concentrations. Therefore, the highest attainable accuracy should be expected

from these measurements as exemplified by the excellent agreement between two recent determinations of HO + C 3 H 6 by Atkinson and Pitts [11]

k = 25.1 ± 2.5 x 10" 12 cm

3 molecule 1 s" 1 and by Ravishankara et al . [12] (k = 25.6 + 1.2 x 10" 12

) at 298 K.

An extensive, critical review of previous

studies on HO-olefin reactions has been made

recently by Atlkinson et al . [10]. In the case of

the lightest olefin, C 2 Hi4 , the pressure dependence

of k has been well established. The rate constant

at 300 K is in the fall -off region between 2nd

and 3rd order kinetics below ^ 225 Torr of Ar and below ^ 300 Torr of He. Table 3 shows a summary

of limiting high pressure data on HO + C 2 H.» taken from the review by Atkinson et al . [10]. There is

a spread of a factor of 2 among these values,

although the recent values are more consistent

but not up to the expected accuracy. The litera- ture values of k for HO + C 3 H 6 are also summarized

in table 3. The fall-off region for C 3 H 6 appears

to occur at much reduced pressure of ^ 1 Torr. Also included in this table are several values derived from relative decay rates of hydrocarbons measured in the photolysis of HC-NO mixtures at ppm concentrations. These relativex values comple- ment those of direct studies, but must be considered to be less precise (by as much as ± 20 percent)

The most extensive and consistent set of both direct and relative rate data for a large variety

of olefins including terpenes and haloalkenes have

been obtained recently by Pitts' group [10]. Their values agree, in general, with those determined by others to within ± 25 percent or better. Therefore, the kinetics of HO-olefin

reactions should be considered to be reasonably

well established. However, it should be noted that

thus far, no direct measurements of these constants

have been made in the presence of 1 atm air. Therefore, it is desirable to obtain further

verification and improvement of these rate constants over wider range of pressures and temperatures.

Mechanism - Primary Step : There have been con- flicting mechanistic interpretation of kinetic and product data obtained at low diluent pressures concerning the relative importance of HO-addition

to olefinic double bond and H-atom abstraction.

However, recent studies show convincingly that HO radicals undergo predominantly, if not exclusively,

addition reactions with C 2 H U , C 3 H 6 and other

methyl -substituted olefins, and that the resulting

adducts are col 1 i sional ly deactivated in 1 atm of

air. For example, in the case of C 2 Hi4, Howard [13]

has shown that the rate constant extrapolates to zero at zero diluent pressure. Thus, the abstraction reaction is of negligible importance. In the case of C 3 H 6 , Cvetanovic [14] has made a comprehensive analysis of products formed in the photolyses

of NzO-CH^-CaHe and N 2 0-H 2 -C 3 H 6 mixtures, and concluded that the majority of the observed

products could be accounted for by reactions of radical species arising from the HO adducts. From these data, Cvetanovic further deduced the extent

of the terminal addition being approximately 65 percent. It is interesting to note that Gutman

et al . [15] earlier obtained evidence for H-abstraction from C 3 H 6 and trans-2-butene under

virtual absence of collisional deactivation in their cross-molecular beam - mass spectrometric

studies. Thus, H-abstraction does occur to a minor extent. A somewhat puzzling study of the H0-C 2 H 4 mechanism is that of Meagher and Heicklen [16]. These investigators photolyzed

H 2 0 2 in the presence of C 2 H 1) and excess N 2 in both the presence and absence of added 0 2 at 298 K. The observed quantum yields of products such as CH 2 0, HCOOH and C 2 H 3 0H were taken to imply that the abtraction accounts for 20 percent of the primary reaction. This conclusion is at odds with those of other workers and must be judged

to be incorrect.
Table 3. Rate constant data and Arrhenius parameters for the reaction of OH radicals with alkenes.'

Al kenes

10 12 x A cm

3 molec 's -ib

E cal mol 10 12 x_kcm

3 molec

_1

s At T K Technique Reference

Temperature

range covered

Ethene

(1 imiting high pressure data)

1.26 -903 ± 136 5.33

L

5.33 6.23

7.21

8.1

0.65u

0.33 0.33

1.6

2.18 -770 ± 300

Propene

4.1

7.85 ± 0.79

10.0 ± 1.7

17 i 4

5.0 ± 1.7 14.5 ± 2.2

13.4 ± 3.4

14.3 ± 0.7

l*kL-ii 20.0 ± 1.0

5 ± 1 -1080 ± 300 25.1 ± 2.5

298-301 300

381 416

305 ± 2

299

296

300

t 300

298

298

381 416

300 298

25.6 ± 1.2 9

27.4 ± 5.5

20.8

23.5 ± 3.5

23.5 ± 4.7

298

FP-KS FP-RF

PR

Relative

rate

FP-RF

FP-RA

DF-MS

DF-ESR

FP-RF

Rel ative

rate

PR

DF-RA

FP-RF

FP-RF

305 ± 2 Relative

rate

303 Rel ative

rate

305 ± 2 Relative

rate

305 ± 2 Relative

rate

Greiner, 1970 (97) a 298-498 K Davis et al . , 1 975 (1 65)

Gordon & Mulac, 1975 (115)

Lloyd et al . , 1 976 (135) (relative to QH+n-butane = 2.82 x 10" 12

) e

Atkinson, Perry & Pitts, 299-425 K 1977 (168) Overend & Paraskevopoulos,

1977 (103)

Morris, Stedman & Niki,

1971 (78) Bradley et al., 1 978 (1 51 ) Stuhl , 1 973 (169) Gorse & Volman, 1 974 (1 23) (relative to QH+CO = 1.50 x 10" 13

) e

Gordon & Mulac, 1 975 ( 1 1 5) Pastrana & Carr, 1975 (170) Atkinson & Pitts, 1975 298-424 K (155) Ravishankara et al . , 1978

(105) Lloyd et al . , 1 976 (1 35) (relative_to QH-n-butane = 2.82 x 10" 12

) e

Wu, Japar & Niki, 1 976 (121 ) (relative to QH+ci s-2-butene = 5.20 x lo" 11 )" Winer et al . , 1976 (136) (relative to OH + isobutene = 4.80 x 10 n ) h Winer et al . , 1977 (140) (relative_to OH + isobutene = 4.80 x lo" n

) n

bTaken from ref. [10]. Reference numbers are indicated as they appear in ref. [10].

, i.e., 100 Torr of he! ium.

305 K.

^Mean Arrhenius preexponential factor

d Total pressure not stated, but stated to be the same as in previous work (86)

e Essentially the high pressure limits from a Lindeman plot (175).

f Calculated from the Arrhenius pressure expression of reference (153) for T = Reference (86).

9 Rate constants at 20 Torr total pressure with helium as the diluent gas. No

h observed over the total pressure range 3-20 Torr (1-butene and cis-2-butene) Calculated from the Arrhenius expressions of reference (155) for T = 303 K or

pressure effects were

or 20-200 Torr (propene).

T = 305 K.

In the case of olefins containing weak allylic

hydrogens, e.g., 1-butene and 3-methyl-l-butene,

Atkinson et al . [17] have postulated, from the correlation of HO and 0(

3 P)-atom reactivities

towards olefins, that H-atom abstraction can occur

up to 30 percent of the total reaction. Clearly, extensive product studies for larger olefins in the presence of an inert diluent gas at 1 atm pressure are needed to obtain more definitive

information on the questions of HO-addition vs. abstraction, and position of the HO-addition.

Secondary Reactions in the Atmosphere : The nature of atmospheric reactions initiated by H0-

olefin reactions is, generally, even less certain

than the corresponding primary processes. However

in the case of G2H4, C 3 H 6 and 2-&,M 8 > for which

the HO-addition has been shown to be the predomin-l ant primary step, some significant progress has been made on the mechanistic interpretation of secondary reactions involving the HO-adducts in the presence of 0 2 and NO . In particular,

computer-aided numerical Analyses of smog chamber

data have played a major role.

ft In computer modeling of smog chamber data, the degree of success in arriving at a unique chemical mechanism may be judged by the extent of agreement

10
0( S P) H02 , N03 , R0X OTHER PRODUCTS (<~2%) (NEGLIGIBLE)

(40%)

[d

1/2 C2 H5CHO

C

2 H5

. 0 + 1/2 C'-C-C (OVERALL ~2%)*

+ HCO

(OVERALL ~l%)»

[fragmentation]

\_ PATH (50%)/ \(50%)

CHjCHO HCHO

+ 2 H02 +CH 3*+H02 + C02 + C02 (OVERALL (OVERALL 3%)** 3%)**

0 OH |ND

II I C-C-C

(OVERALL

21%)""

CH3 CHO + HCHO

(FROM 03 + c=c-c: OVERALL 24%)**

CHjCHO + HCHO L°_ (FROM OH + C =C-C : OVERALL '- 43%)**?

Fig. 2. Atmospheric reactions involving C 3 H 6 (from ref. 19).

between computed and experimental concentrationtime profiles of all the relevant reactants and

products. Unfortunately, many existing smog chamber data have not been designed for the purpose of elucidating detailed reaction mechanisms, and consequently, suffer from the lack of sufficiently accurate and extensive product

analyses and other kinetic parameters. In addition, for typical smog chamber data on olefins,

both HO- and 0 3 - initiated oxidation of olefin

and No are of comparable importance. Thus, mechanistic understanding of HO-olefin reactions

requires adequate knowledge of the corresponding

0 3 -olefin reactions or vice versa. Further complication in the interpretation of smog chamber

data arises from the highly ill-defined, elusive

role of what is commonly known as "chamber

effects" [18]. These effects appear to be significant source and sink for chain carriers,

e.g. , HO and H0 2 . It is apparent that better characterization of such effects is a prerequisite

to the judicious use of smog chamber data for validating detailed smog mechanisms.

The current state-of-the-art smog chamber data generation and computer modeling can be illustrated

by the work of Pitts' group [19] with CsHs-C^Hio

mixtures as surrogate hydrocarbons. The H0-C 3 H 6 mechanism adopted by these investigators is shown

in figure 2, together with other atmospheric

reactions involving C 3 H 6 . In this scheme, two of

the most important radical intermediates in the

OH oxidation of the terminal HO adduct, (C-C-C), are

OH 0

I I C-CC and CH 2 OH . The corresponding radicals for

OH O-OH

the non-terminal HO adduct, (C-C-C), are C-C-C

and CH 3 CH0H. The former oxy-type radicals undergo predominantly unimolecular dissociation, and

to a minor extent, bimolecular reactions with 0 2 and N0 2 . The reactions of the latter hydroxyalkyl

radicals with 0 2 lead to H-atom transfer from the HO-group rather than the addition of 0 2 to form peroxy radicals. Presumably, this mechanism has been deduced from studies of the NO -photoxidations

of ethanol and 2-butanol . The same

x concl usion

has been reached independently by Niki et al . [20] from Fourier Transform Infrared (FTIR) studies of Cl-atom initiated oxidation of NO alcohol mixtures

in ppm concentration.

According to the scheme in figure 2, the HO- initiated oxidation of C 3 H 6 in the presence of NO leads to the formation of HCHO and CH 3 CH0 as the major carbon-containing products and to the concomitant conversion of 2N0 -> 2N0 2 .

Similarly, the major products from CzH^ and 2-C„He are expected to be HCHO and CH 3 CH0, respectively. Further evidence to support this conclusion has been recently obtained by Niki et al . [20] from product studies in the photo- oxidation of these olefins in the presence of HONO

and NO. These experiments have been designed to produce HO from a well -defined photochemical

source, i.e., HONO + hv -» HO + NO, with minimal interference from 0 3 -olefin reactions by maintaining a large excess of NO over N0 2 . Product analyses by the FTIR were made at sufficiently

low conversion of the olefins to minimize

secondary reactions of HO with the aldehydic

products. Thus, this work represents a significant

11
departure from conventional smog chamber-based

studies, and should be extended to other olefinic

compounds

.

In conclusion, there is a great need for kinetic

and mechanistic data on the oxidation of free radicals formed from HO-olefin reactions. The above-mentioned studies illustrate a classical photochemical approach to these problems. Clearly,

it is highly desirable to study these chemical systems with more direct experimental methods.

B. 0 3 -olefin reactions

The atmospheric consumption of olefins proceeds,

to a large extent, by their reactions with 0 3

. Furthermore, the question of whether the 0 3 -olefin reactions serve as an 0 3 source or sink hinges on the efficiency of free radical formation by the reaction. Thus, utmost accuracy is required for

the kinetic and mechanistic information on 03- olefin reactions for modeling purpose.

Kinetics : Over the years, there have been numerous determinations of the rate constants for

the reactions of 0 3 with a large variety of olefins. A summary of literature values for several olefins is given in table 4. Generally,

the reported values agree reasonably well for

terminal olefins, but scatter far beyond the estimated experimental precision for internally double-bonded olefins. For instance, two of the most recent sets of extensive measurements by Huie and Herron [21] (ref. [1] in table 4), and

by Niki et al . [8] (refs. f and k) disagree by as much as 48 percent for trans-2-butene and by an average of 27 percent for the entire set. The individual values in the latter set are all higher

than those in the former. Apparently, there are factors other than systematic measurement errors affecting one or both experiments. In these studies as well as most of the others, the rate constants were derived from the decay rates of 0 3 in the presence of excess olefins. In addition,

the former employed much higher reactant concentrations and lower total pressures (< 10 Torr) than the latter (at 1 atm air). Therefore, the fundamental question is whether the consumption of

0 3 is due entirely to the primary process or is interfered with by secondary reactions. Unfortunately, there exist large uncertainties in the reaction mechanisms, and such effects cannot be determined reliably at present.

There are several instances of kinetic evidence

for the consumption of 0 3 by secondary reactions,

e.g., retardation of excess 0 3 decay rates by

0 2 , [21, 22] and dependence of 0 3 and olefin decay

rates and reaction stoichiometry on reactant mixing ratios [22]. Therefore, all the reported

values of the rate constants for 0 3 -olefin reactions must be considered as upper limit values.

In particular, the higher values obtained by Niki et al . [8] might reflect the extent of such effect.

In short, experimental methodologies required

for obtaining the "true" rate constants for 03- olefin reactions have not been firmly established

at present, and should be developed in the future. Meanwhile, concerted efforts should be made to minimize systematic errors as revealed in table 4. Concomitantly, kinetic and mechanistic studies of

Table 4. Rate constants for qas phase ozone-olefin

reactions at room temperature. 3

k , 1 0~ 1

8

cm

3 molec"^" 1 Literature

Ethylene

Propyl ene

1 -Butene

1 -Pentene

1-Hexene

Dial kylethylenes

Isobutene

ci s-2-Butene

trans-2-Butene

Trial Methylenes

2-Methyl-2-butene

ci s-3-Methyl -2- pentene trans-3-Methyl -2- pentene

Tetraal kyl ethylenes 2,3-Dimethyl-2-

butene

Others Cyclopentene

Cyclohexene

1 ,3-Butadiene

1.2

e

2.7

d 6.2, 7.5

L

- f 13

1

1 .3

U

, 1.6' , 2.S

3.0

3

, 1.9

k

, 1.7

1

B.2 , 11.0

12.5

9.0' 5.3

a

, 7.5

b

, 9.0C

10.6'

10.

0

d

, 12. 3

k

, 10. 3

1

10.7'

; d ,f,h i n.o

6.2 , 8.4 L

11.1'

11.7

h , 13.6

28

C

, 50' 126

1

35

C

, 166

h

, 260

1

, 275

f

, 430 d

, 260

k

, 176

1

29

c

, 790

1

, 493 , 400

456

k

563

k

39

C

, 750

d

, 1510

k

, 1060

1

813

30

1

, 59

a

, 169

1"

8.2

b

, 9.1

1

, 8.

Taken from ref. [8]. References (a) through (k) given in ref. [8]; reference (1) is ref. [23] of

this text. free-radical reactions involving 0 3 and olefins should be made to better characterize the relevan secondary reactions as discussed below.

Mechanism : To date, no direct monitoring of reactive intermediates formed in 0 3 -olefin reactions has been made under atmospheric

conditions. Therefore, it is compelling to resor, to the deduction of mechanistic models from produ studies. The model will provide a basis for establishing experimental priorities to reduce mechanistic uncertainties. Significant progress

has been made recently both in characterizing key reaction products and in computer modeling of these results.

To illustrate, the reaction of 0 3 with C2H., ha been studies by Herron and Huie [23], using the stopped-flow mass spectrometry (M.S.) method, at 298 K and 8 Torr total pressure. From the comput

modeling of the temporal behaviors of several species, e.g., Z 2 W W , C0 2 , H 2 0, CH 2 0, HC00H and

CH 3 0H, the role of free radical mechanism involvi

27 steps has been proposed.

12
i The initial reactions occurring under their

Isxperimental conditions have been postulated to be

(Criegee mechanism [29] followed by the unimolecular

dissociation of the methylene peroxide, CH 2 00, to

serveral products, i.e.,

CH 2 0 2

-> H 2 C00 67%

18%

9%

6%

H 2 0 + CO

H 2 + C0 2

2H + CO2

HCOOH

I The formation of intermediate dioxirane (H 2 C00)

iin the above scheme was predicted by Wadt and

ioddard [24], and was subsequently verified by

.ovas and Sueram [25] using microwave spectroscopy

in the low-temperature reaction of 0 3 with C 2 Hi,.

Clearly, the degree of reliability of the above

>ostulated mechanism depends on the accuracy of the

nput kinetic data for the series of secondary

eactions initiated by the product H atoms. On the

pasis of their modeling work, Herron and Huie [23]

;iave pointed out the needs for improved data for

everal crucial free radical reactions occurring

n Q 3 -C 2 Hk system. In particular, the mechanistic

nowledge of the H0-C 2 hU reaction is a prerequisite

0 unraveling the 0 3 -C 2 Hi» mechanism. These nvestigators further extended their model to high

iressure conditions and compared the computed

Results with the product data on Q 3 -C 2 Hh reactions

bbtained by Scott et al . [26] at ppm reactant oncentrations in 1 atm air. In general, the

igreement should be considered fair considering

he uncertainty in both measurements and model.

Very recently, Dodge and Arnts [27] have eveloped a sufficiently detailed model including

6 steps for the reactions of 0 3 with methylubstituted olefins, e.g., C 3 H 6 and 2-C^H 8 , which

nvolve the formation and subsequent dissociation

f a "hot" CH 3 CHOO radical, i.e.,

P3.

C-00

+ 20%

CH 3 HCHO X

C — 00 Ozonide

80% CH 3 ^0

nd

The above scheme for the decomposition of the

CH 3 CH00 radical to several free radical species was used to model the product studies of Niki et al . [28] in the 0 3 -air-2 butene - HCHO - air system.

Although these recent experimental and modeling

studies are by no means definitive, they do shed new light on the 0 3 -olefin chemistry, in particular, on the efficiency of radical formation. Some of the 0 3 -olefin mechanisms adapted for atmospheric modeling incorporate the formation of two free radical species for every olefin molecule

reacted with 0 3 [9]. This assumption stems from

two types of mechanistic conjectures which appeared

plausible in the past. One of these was based on the motion that secondary ozonides are not formed from small olefins in air because of chemical instability of the corresponding Criegee intermediates, possibly reacting with 02 [26]. The recent observation of propene ozonide in a mixture of

0 3 -2-C^H e -HCH0-air precludes this possibility [28]. The other stems from a theoretical treatment of

gas phase 0 3 -olefin reactions by O'Neal and Blumstein [30] which deviates from the Criegee

mechanism. At present, there is little experimental evidence that supports their theory uniquely

Thus, the new information suggests the production

of fewer radical species than has been assumed

previously.

4. Conclusion

Experimental methodologies for the determination

of the kinetics of HO-olefin reactions appear to be well established, and should be applied more extensively to numerous olefins under wider range of temperature and diluent pressure. In contrast, derivation of "true" rate constants for 0 3 -olefin reactions from the decay rates of the reactants may suffer from the interferences due to secondary free

radical reactions initiated mainly by "Criegee

intermediates," and thus requires better mechanistic knowledge. To elucidate the mechanisms

for 0 3 -olefin reactions, temporal behavior of the reactants and products should be better characteriz

ed first. These data can then be utilized for mechanistic modeling. The HO-initiated oxidation

of olefin is likely to play a major role in the secondary reactions of 0 3 -olefin system. Thus,

mechanistic studies of HO-olefin reactions in the presence of 1 atm air are prerequisite to the understanding of 0 3 -olefin reactions, and should assume the highest research priority.

CH 3 /0

Xyt CH k + C0 2 /

C J H 0 68%

0

H COCH3

CH 3 + CO + HO

\J

0%

CH 3 + C0 2 + H

50%

HC0 + CH 3 0

\!j 0%

H + C0 2 + CH 3

References

[1] Demerjian, K. L. , Kerr, J. A., and Calvert,

J. C, Adv. Environ. Sci . Technol . 4, 1 (1974).

[2] Dodge, M. C. , Effect of Selected Parameters

on Predictions of a Photochemical Model,

EPA-600/3-77/148 (June, 1977).

[3] Echenroeder, A., Development of an Air Quality Model Based on the LARPP Data, (in progress)

[4] Parker, R. 0. and Martinez, J. R. , Los Angeles Reactive Pollutant Program (LARPP) Data Archiving and Retrieval, Document No.

13
P-1464-W, 1974, Environmental Research and Technology, Inc. Concord, Massachusetts,

(NTIS Accession No. PB 244-298).

[5] Garvin, D. and Hampson, R. F. eds., Chemical Kinetics Data Survey VII, Table of Rate and Photochemical Data for Modelling of the Stratosphere [Revised], NBSIR 74-430,

National Bureau of Standards, Washington,

D.C. 20234 (May, 1973).

[6] Leighton, P. A., Photochemistry of Air

Pol 1 ution (Academic Press, New York, 1961).

[7] Calvert, J. G. , Environ. Sci. Technol . 10, (1976).

[8] Japar, S. M. , Wu, C. H. , and Niki, H.,

J. Phys. Chem . 78, 2318 (1974).

[9] Niki, H., Daby, E. E. , and Weinstock, B., Adv. Chem. Ser . 113 , 16 (1972).

[10] Atkinson, R. , Darnall, K. R., Lloyd, A. C,

Winer, A. M. , and Pitts, J. N. , Jr., Adv. Photochem . (in press).

[11] Atkinson, R. and Pitts, J. N. , Jr., J. Chem . Phys . 63, 3591 (1975).

[12] Ravishankara , A. R. , Wagner, S., Fischer, S., Smith, G., Schiff, R. , Watson, R. T., Tesi

,

G. , and Davis, D. D., Int. J. Chem. Kinet . 10, 783 (1978).

[13] Howard, C. J., J. Chem. Phys ., 65 4771 (1976).

[14] Cvetanovic, R. J., 12th Int. Symp. Free Radicals, Laguna Beach, California, Jan. 4-

9 (1976).

[15] Slagle, I. R. , Gilbert, J. R. , Graham, R. E., and Gutman, D., Int. J. Chem. Kinetics , Symp. No. 1 , 317 (1975).

[16] Meagher, J. F. and Heicklen, J., J. Phys. Chem . 80, 1645 (1976).

[17] Atkinson, R. , Perry, R. A., and Pitts, J. N.

,

Jr., J. Chem. Phys . 67_, 31 70 (1977).

[18] Bufalini, J., Walter, T. , and Bufalini, M., Environ. Sci. Technol . V\_, 1 181 (1977).

[19] Carter, W. P. L., Lloyd, A. C. , Sprung, J.

R. , and Pitts, J. N. , Jr., Computer Modeling

of Smog Chamber Data: Progress in Validation

of a Detailed Mechanism for the Photooxidation of Propene and p-Butane in Photochemical

Smog, Int. J. Chem. Kinetics 11 , 45 (1979).

[20] Niki, H., Marker, P. D. , Savage, C. M. , and Breitenbach, L. P., J. Phys. Chem . 82, 135 (1978).

[21] Huie, R. E. and Herron, J. T. , Int. J. Chem . Kinetics , Symp. 1 , 165 (1975).

[22] Japar, S. M. , Wu, C. H., and Niki, H., J^_ Phys. Chem . 80, 2057 (1976).

[23] Herron, J. T. and Huie, R. E., J. Am. Chem . Soc . 99, 5430 (1977).

[24] Wadt, W. R. and Goddard, W. A., Ill, J. Am . Chem. Soc . 97, 3004 (1975).

[25] Lovas, F. J. and Suenram, R. 0., C hem. Phys , Lett . 51, 453 (1977).

[26] Scott, W. E., Stephens, E. R. , Hanst, P. L. ; and Doerr, R. C, Proc. Am. Petroleum Inst . 37., 1 71 (1957).

[27] Dodge, M. C. and Arnts, R. R. , A New Mechanism for the Reaction of Ozone with

Olefins, Int. J. Chem. Kinetics 11, 399

(1979).

[28] Niki, H., Maker, P. D. , Savage, C. M. , and Breitenbach, L. P., Chem. Phys. Lett . 46,

327 (1977).

[29] Criegee, R. , Rec. Chem. Progs . 18, 113

(1975).

[30] O'Neal, H. E. and Blumstein, C, Int. J . Chem. Kinetics 5, 397 (1973).

Summary of Session

The discussion centered on the reactivity of

OH and 0 3 and the mechanisms of reaction of OH and' 0 3 with olefins. The mechanisms of OH-olefin

reactions were treated by Cvetanovic. In his work hydroxyl radicals were produced by the photolysis of N 2 0 (to produce 0'D) in the presence of

H 2 or H 2 0. In the absence of 0 2 , the major

products of the reaction can be accounted for on the basis of an additions reaction followed by radical -radical and radical substrate reactions. There was no evidence that abstraction was important for ethylene or propylene, e.g., in the case of the propylene reaction only minor amounts

of products characteristic of vinyl radicals, its major abstraction product, were found. Heicklen discussed the earlier work which had suggested the possible importance of an abstraction path for

OH + C 2 H h . As pointed out by Niki in his review

paper, larger olefins, particularly those with weak allylic bonds, all could react partially

through an abstraction mechanism. Kinetic data support this argument.

Of greater interest is the question of the rol of 0 2 in the reaction. If the initial reaction i addition, then in the presence of 0 2 , a peroxy

radical would be formed. The subsequent fate of this peroxy radical is one of the key problems

in laboratory studies of atmospheric chemical reactions. Cvetanovic reported that, in the presence of 0 2 , the major products of the

OH + C 2 H„ reaction were CH 3 CH0 and C 2 H 5 0H, and from OH + C 3 H 6 , the major product by far was

CH 3 CH0. These products are not readily accounted

for on the basis of simple radical -radical interactions 2R0 2 R'CHO + R"0H + 0 2 , or 2R0 2

-> 2R0 + 0 2

. In the atmosphere where NO is

14
1 available it is generally assumed that the peroxy

| radicals are converted to alkoxy radicals. The

I subsequent fate of the alkoxy radicals is not understood. In the mechanism used by Niki (his

; fig. 2) based on modeling of the Riverside data,

j the alkoxy radical decomposes (a revised version

of the mechanism and figure were supplied by

Carter). It was pointed out by Batt and by

Golden, however, that the simpler alkoxy radicals

were more likely to react with O2 (see review

I paper of Golden and susequent discussion). The

j specific fate of the radical C 3 H 6 -0H'02 produced

1 in the reaction of OH with C 3 H 6 in the presence

j of 0 2 was not resol ved.

The rate constants for OH reactions are in relatively good shape as discussed by Atkinson.

Rate constants for 0 3 reactions are still subject

to some uncertainty because of problems in measurement methodology. These problems would be resolved if the mechanisms were better understood,

j The major lack of kinetic data is for cyclic

(olefins which as Niki pointed out are the most

I important class with respect to consumption of

I ozone by olefins.

The mechanisms of ozone-olefin reactions have

i not been settled. Recent work, as reviewed by

Niki, indicates that free radical yields are smaller than had been thought previously. Whitten

emphasized that in most cases the radical yields

J| were the important information needed by modelers

| rather than specific reaction channels. Dodge

|] presented a general model for the reactions based

I on both high and low pressure experimental observaI tions. This model appears to rule out the O'Neal

Blumstein mechanism. O'Neal agreed to this interpretation, pointing out the reasons for the

discrepancy. Golden pointed out that the high

and low pressure results used by Dodge might not

J be compatible because of different quenching rates

I etc., and the general mechanism should be used with

caution. A critical question in this discussion

is the chemistry of the Criegee intermediate,

RCHOO. In particular what are the rates of

reaction with NO , SO , and aldehydes relative

to i somerization xand Scission. Heicklen commented

on the reactions of ozone with chlorinated

ethylenes where a ir-complex may be involved.

There was considerable interest expressed in extending observations to more complex systems.

The importance of cyclozlkenes was noted above.

iStedman urged consideration of natural substances

1 such as a-pinene and isoprene. O'Brien pointed

.out the difference between high and low molecular

weight compounds particularly with respect to acid and aerosol formation. This raise unanswered

questions as to the stabilization of the Criegee

intermediate and their role in the formation of

acids. The origin of organic acids is also a puzzling problem in photooxidation studies of aldehydes. Acids are also found in low pressure

jozonoalysis studies.

Comments

Roger Atkinson, Statewide Air Pollution Research

Center, University of California, Riverside, California 92521 Rate Constant Data for the Reaction of OH Radicals with Alkenes : The available rate constant

data and Arrhenius parameters are given in tables

1 (alkenes), 2 (monoterpenes ) and 3 (dialkenes) which are taken from the review article of Atkinson, Darnall, Lloyd, Winer and Pitts, to be published in Advances in Photochemistry . The relative rate~constant data of Cox [1] for ethane, propene and trans-2-butene, obtained from

the photolysis of HCHO-alkene mixtures in air at 760 Torr total pressure, have not been included in table 1 as the stoichimetry factor was not known. The data from relative rate studies have been reevaluated on the basis of what is felt (however

biased) by the authors to be the "best" rate constant for the reference reaction at the temperatures employed in the respective relative

rate studies. The abbreviations used for the techniques are: DF-RF: discharge flow-resonance

fluorescence; DR-RA: discharge flow-resonance

absorption; DF-ESR: discharge flow-esr detection;

DF-LMR: discharge flow with laser magnetic

resonance detection of OH; DF-MS: discharge flowmass spectrometry; FP-KS flash photolysis-kinetic

spectroscopy; FP-RA: flash photolysis-resonance

absorption; FP-RF: flash photolysis-resonance

fluorescence; MPS: modulation-phase shift; PR: pulsed radiolysis. The tables list the literature

data expressed in the form of k = A e~E/RT with k being given at room temperature, and A and E are

the Arrhenius preexponential factor and Arrhenius

activation energy, respectively.

For ethane the rate constant at room temperature (and up to 425 K) is in the falloff region between second order and third order kinetics [3,6,7,27,31] below ^ 225 Torr of argon [6] and below ^ 300 Torr of helium [3]. The rate in table

1 for ethene are limited to the high pressure re- sults, although the rate constants of Greiner [2] (apparently obtained at 100 Torr total pressure of

helium) have been included.

For ethane and propene the limiting high pressure room temperature rate constants appear to

be reasonably well defined (table 1) at about 8 x 10" 12 cm 3 molec -1

s _1 and 2.5 x 10" 11 cm 3 molec -1

s _1

, respectively. For the higher alkenes the data do not appear to be as consistent. The rate constants obtained by Ravishankara et al . [18] for cis-2-butene and, especially, for 2,3-dimethyl-2-

butene appear to be low; that for 2,3-dimethyl-2-

butene [18] being a factor of 2 lower than the room temperature rate constants determined by Morris and Niki [19], Perry [25], and Atkinson,

Darnall and Pitts [23]. It is probable that wall adsorption problem (as observed for propene in a metal reaction cell L18]) in the static system

used by Ravishankara et al . [18] is the cause of these apparently low rate constants.

It would then appear that the flash photolysisresonance fluorescence rate constant data of Atkinson and Pitts [14,22] and Atkinson, Perry,

and Pitts [6,20] (which are generally in good

15
Table 1. Rate constant data and Arrhenius parameters for the reaction of OH radicals with alkenes.

10 12 x A

Al kenes

10 12 x_k

E cal mol" 1 cm

3 mo1 ec" 1 s" 1 At T K Technique Reference

Temperatu

range covered

Ethene

(limiting high pressure data)

1.26

2.18

-903 ±136 5.33"

5.33 ± 0.65c

6.23 ± 0.33

7.31 ± 0.33

8.1 ± 1.6

-770 ± 300 7.85 ± 0.79

10.0 ± 1.7

298-301 FP-KS

300 FP-RF

381 416)

PR

305 ± 2 Relative

rate

299 FP-RF

296 FP-RA

Ethene-d.,

Propene

8.24 ± 0.48 298 ± 2 Relative

rate

4.1

17 ± 4 300

5.0 ±1.7 ^300

14.5 ± 2.2 298

13.4 ± 3.4 298

14.3 ± 0.7 381 20.0 ± 1.0 416

5 ± 1 300

-1080 ± 300 25.1 ± 2.5 298

27.4 ± 5.5

DF-MS

DF-ESR

FP-RF Relative

rate

PR

DF-RA

FP-RF

305 ± 2 Relative

rate

20.8 303 Relative

rate

23.5 ± 3.5

23.5 ± 4.7

305 ± 2 Relative

rate

305 ± 2 Relative

rate

Propene-d 6

1-Butene

7.6

25.6 ± 1.2" 298 FP-RF

18.7 298 DF-MS

16.8 298 FP-RF

40.8 298 DF-MS

15 ± 1 300 DF-RA

-903 ± 300 35.3 ± 3.6 298 FP-RF

27.0 303 Relative

rate

29.4 ± 1.4

h 298 FP-RF

Greiner, 1970 [2] 298-498 I Davis et al. , 1975 [3]

Gordon & Mulac, 1975 [4]

Lloyd et al . , 1976 [5] (relative to QH+n-butane = 2.82 x 10~ l2

) a

Atkinson, Perry & Pitts, 299-425

1977 [6] Overend & Paraskevopoulos,

1977 [7]

Niki et al . , 1978 [8] (relative to QH + C 2 H^ 8.00 x 10" 12

) e

Morris, Stedman & Niki,

1971 [9}

Bradley et al. , 1973 [10] Stuhl, 1973 [11] Gorse & Volman, 1974 [12] (relative to OH+OC = 1.50 x 10~ 13 )f

Gordon & Mulac, 1975 [4] Pastrana & Carr, 1975 [13] Atkinson & Pitts, 1975 298-424 K [14] Lloyd et al. , 1976 [5] (relative to OH-n-butane = 2.82 x 10" 12 )d Wu, Japar & Niki, 1976 [15] (relative to QH+cis-2-butene = 5.20 x 10" 11

) 9 Winer et al. , 1976 [16] (relative to QH + isobutene = 4.80 x 10" 11

) 9 Winer et al . , 1977 [17] (relative to QH + isobutene = 4.80 x 10" 11

) 9 Ravishankara et al., 1978

[18]

Morris & Niki, 1971 [19] Stuhl, 1973 [11]

Morris & Niki, 1971 [19] Pastrana & Carr, 1975 [13] Atkinson & Pitts, 1975 298-424 K [14]

Wu, Japar & Niki , 1976

[15] (relative to OH + cis-2-butene = 5.20 x 10" 11

) 9 Ravishankara et al., 1978

[18]

16
Table 1. Rate constant data and Arrhenius parameters for the reaction of OH radicals with alkenes. (continued)

Al kenes

10 12 x A 10 12 x k cm

3 molec

_1 s~

ia

E cal mol" 1 ctn

3 molec

_1

s -1 At T K Technique Reference

Temperature

range covered

Isobutene

cis-2-Butene

trans-2-Butene

1-Pentene

P-Methyl- 1-butene

3-Methyl2-Methyl- 2-butene

:is-2- Pentene

9.2

10.4

11.2

5.23

36

19.1

64.6

-1000 ± 300 50.7 ± 5.1

47.8

61.2

-970 ± 300 53.7 ± 5.4

298 DF-MS

297 FP-RF

303 Relative

rate

298 DF-MS

298 FP-RF

61.5 ± 12.3 305 ± 2 Relative

rate

58.6 ± 8.8 305 ± 2 Relative

rate

42.6 ± 2.5'

71.4

12 ± 10

-1090 ± 300 69.9 ± 7.0

67.6

42.5

29.1

90.1 57.2

-1060 ± 300 31.0 ± 3.1

119 -450 ± 400 78

-895 ± 300 87.3 ± 8.8

87 ± 6

62.4

298 FP-RF

298

300

298

303

298

303

298

303

DF-MS

DF-RA

FP-RF

Relative

rate

DF-MS

Relative

rate

DF-MS Relative

rate

299 FP-RF

298 DF-MS

298 FP-RF

299 FP-RF

300 ± 1 Relative

rate

303 Relative

rate

Morris & Niki, 1971 [19] Atkinson & Pitts, 1975 297-424 K [14]

Wu, Japar & Niki , 1976 [15] (relative to OH + cis-2-

butene = 5.20 x 10" 11

) 9

Morris & Niki , 1971 [19] Atkinson & Pitts, 1975 298-425 K [14] Lloyd et al. , 1976 [5] (relative to OH + n- . butane = 2.82 x 10" 12

) a

Winer et al . , 1976 [16] (relative to OH + isobutene = 4.80 x 10" 11

) 9

Ravishankara et al . , 1978

[18]

Morris & Niki, 1971 [19] Pastrana & Carr, 1975 [13] Atkinson & Pitts, 1975 298-425 K [14]

Wu, Japar & Niki , 1976

[15] (relative to OH + cis-2-butene = 5.20 x 10" 11

) 9

Morris & Niki, 1971 [19]

Wu, Japar & Niki , 1976

[15] (relative to OH + cis-2-butene = 5.20 x 10" 11

) 9

Morris & Niki , 1971 [19] Wu, Japar & Niki, 1976 [15] (relative to OH + cis-2- butene = 5.20 x 10" 11

) 9

Atkinson, Perry & Pitts, 299-423 K 1977 [20]

Morris & Niki , 1971 [19] Atkinson, Perry & Pitts, 298-425 K 1976 [21] Atkinson & Pitts, 1978 299-426 K [22] Atkinson, Darnall & Pitts,

1978 [23] (relative to

OH + cis-2-butene = 5.29 x 10" 11

) 9

Wu, Japar & Niki , 1976

[15] (relative to OH + cis-2- butene = 5.20 x 10" 11

) 9

17
Table 1.

Al kenes

2-Pentene 90.1 298 DF-MS Morris & Niki, 1971 [19]

(mix. cis,

trans)

1-Hexene 31.2 303 Relative

rate

Wu, Japar & Niki, 1976

[15] (relative to OH + cis- 2-butene = 5.20 x 10" 11

) 9

Cyclohexene 62.4 303 Relative

rate

Wu, Japar & Niki , 1976

[15] (relative to OH + cis-2-butene = 5.20 x TO" 11

) 9

-- 73.4 + 14.7 305 ± 2 Relative

rate Darnall et al . , 1976 [24]

( tgI flti vfi to OH + i so~ butene = 4.80 x 10" 11

) 9

3,3-Dimethyl1-butene27.0 303 Relative

rate

Wu, Japar & Niki, 1976

[15] (relative to OH + cis-2-butene = 5.20 x 10" 11

) 9

2,3-Dimethyl- 153 298 DF-MS Morris & Niki, 1971 [19] 2-butene 110 + 22 298 FP-RF Perry, 1977 [25]

56.9 + 1.3

h 298 FP-RF Ravishankara et al . , 1978

[18]

-- 122 ± 8 300 ± 1 Relative

rate Atkinson, Darnall & Pitts, 1978

[23] (relative to OH + cis-2-

butene = 5.29 x 10" 11

) 9

1-Heptene 35.0 ± 7.0 305 ± 2 Relative

rate Darnall et al . , 1976 [24] (relative to OH + isobutene = 4.80 x 10" 11

) 9

1 -Methyl - cyclohexene

91.7 ± 18.3 305 ± 2 Relative

rate Darnall et al . , 1976 [24] (relative to QH + isobutene = 4.80 x 10" 11

) 9

.Mean Arrhenius preexponential factor.

Total pressure not stated, but stated to be the same as in previous work [26], i.e., 100 Torr of helium. .Essentially the high pressure limit from a Lindemann plot [27]. Calculated from the Arrhenius expression of reference [28] for T = 305 K. ^Calculated from the Arrhenius expression of reference [6] for T = 298 K. Rate constant determined for OH + CO at room temperature and low pressure [29]. ^Calculated from the Arrhenius expression of reference [14] for T = 300, 303 or 305 K. Rate constants at 20 Torr total pressure with helium as the diluent gas. No pressure effects were observed over the total pressure range 3-20 Torr (1-butene and cis-2-butene) or 20-200 Torr (propene).

In addition to the above data, Simonaitis and Heicklen [30] obtained rate constant data for propene relative to those for the reaction of OH radicals with CO at 373 and 473 K at total pressures of ^ 400- 800 Torr (mainly H 2 0). They obtained k(0H + C 3 H 6 )/k(0H + CO) (± 10%) = 75 at 373 K and 55 at 473 K. Assuming that k(0H + CO) = 3.0 x 10~ 13 cm 3 molec"*s _1

, independent of temperature under these conditions

(which is subject to large uncertainties), then by extrapolation a value of k(0H + C 3 H 6 ) = 3.3 « 10" 11 cm 3 molec^s" 1 at 298 K and an Arrhenius activation energy of E = -1090 cal mol" 1 may be obtained. It

is evident that this data (which is obviously subject to large uncertainties because of the assumptions

made) is in general agreement with that obtained by references [5.14-18J, as quoted in the table above.

18
Table 2. Rate constant data for the reaction of OH radicals with monoterpenes

10 11 x k Terpene cm 3 molec" 1 s" 1 At T K Technique Re ference

a-Pinene

g-Pinene

5.5 ± 0.

5.4 ± 1.0

d-Limonene 14.2 ± 2.1

305 ± 2 Relative Winer et al . , 1976 [16] rate (relative to OH + isobutene

= 4.80 x 10

_11 ^ a

305 ± 2 Relative Winer et al . , 1976 [12] rate (relative to OH + isobutene

= 4.80 « 10"- 1 1 \3

305 ± 2 Relative Winer et al . , 1976 [16] rate (relative to OH + isobutene

= 4.80 x 10"• 1 1 \S

Calculated from the Arrhenius expression of reference [14] for T = 305 K.

Table 3. Rate constant data and Arrhenius parameters for the reaction of OH radicals with dialkenes.

TO 12 x A

Dialkenes E cal mol 10 12 x_k

_ cn^molec^s" 1 At T K Technique Reference

Temperature

range covered

Propadiene

1 ,3-Butadiene

5.59

14.5

4.5 ± 2.5

U ^ 300 DF-ESR

-305 ± 300 9.30 ± 0.93 299 FP-RF

Bradley et al . , 1973

[10] Atkinson, Perry, & Pitts, 1977 [20]

72.8 + 14.6 305 ± 2 Relative Lloyd et al . , 1976 [5] rate (relative to OH + n-

-930 ± 300 68.5 ± 6.9 299 FP-RF

butane = 2.82 x 10" 12

) c

Atkinson, Perry, & Pitts, 1977 [20]

299-421 K

299-424 K

^Mean Arrhenius preexponential factor.

c May be in the fall-off region between second order and third order kinetics (see text and reference [20]). Calculated from the Arrhenius expression of reference [30] at 305 K.

greement with the relative rate data [5,15-17,23]

nd with the discharge flow-mass spectrometric

lata of Morris and Niki [19]), together with the

ate constant for 2,3-dimethyl-2-butene recently

btained by Atkinson, Darnall, and Pitts [23]

rom a relative rate study, should be viewed as

.he most consistent set of rate constant data,

his is especially so as this set of absolute rate

ionstant data [14,20,22] also comprises the only

[emperature dependence studies for the alkenes

fther than for ethene.

1 Finally, it should also be noted that with the

lash photolysis systems used to determine OH

adical rate constants for the alkenes, problems

ave been encountered due to secondary reactions

nd due to wall absorption of the reactants.

hus, although the [reactant]/[OH] ratios are easonably similar with FP-RA and FP-RF, because

f the higher flash energies used with RA etection 1000 joules per flash compared with

i 100 joules per flash for RF detection), secondfry reactions of OH radicals with the larger

mounts of photolysis products generated by the

jore intense flash may be encoutered with flash

Ihotolysis-resonance absorption systems.

References

[1] Cox, R. A., Int. J. Chem. Kinet . Symp. No. 1, 378 (1975).

[2] Greiner, N. R. , J. Chem. Phys. 53, 1284 (1970).

[3] Davis, D. D. , Fischer, S., Schiff, R., Watson,

R. T., and Bollinger, W. , J. Chem. Phys . 63, 1707 (1975).

[4] Gordon, S. and Mulac, W. A., Int. J. Chem . Kinet. , Symp. No. 1, 289 (1975).

[5] Lloyd, A. C, Darnall, K. R., Winer, A. M.

Pitts, J. N . , Jr., J. Phys. Chem. 80 and Pitts, J. N . , Jr., J. Phys. Chem . 80,

789 (1976)

[6] Atkinson, R., Perry, R. A., and Pitts, J. N .

Jr., J. Chem. Phys . 66, 1197 (1977).

[7] Overend, R. and Paraskevopoul os , G., J. Chem . Phys. 67, 674 (1977).

[8] Niki, H . , Maker, P. D., Savage, C. M., and Breitenbach, L. P., J. Phys. Chem . 82_, 132 (1978).

19
[9] Morris, E. D., Jr., Stedman, D. H., and

Niki, H., J. Amer. Chem. Soc . 93, 3570 (1971).

"10] Bradley, J. N., Hack, W., Hoyermann, K., and Wagner, H. Gg., J. Chem. Soc. Faraday Trans. I 69, 1889 (1973).

[11] Stuhl , F., Ber. Bunsenges, Phys. Chem . 77 , 674 (1973).

[12] Gorse, R. A. and Volmam, D. H. , J. Photochem . 3. 115 (1974).

[13] Pastrana, A. V. and Carr, R. W., Jr., J. Phys . Chem . 79, 765 (1975).

[14] Atkinson, R. and Pitts, J.N., Jr., J. Chem . Phys. 63, 3591 (1975).

[15] Wu, C. H. , Japar, S. M. , and Niki, H., J^

Environ. Sci. Health All , 191 (1976).

[16] Winer, A. M. , Lloyd, A. C, Darnall , K. R., and Pitts, J. N . , Jr., J. Phys.. Chem . 80, 1635 (1976).

[17] Winer, A. M. , Lloyd, A. C, Darnall, K. R., Atkinson, R., and Pitts, J. N . , Jr., Chem.

Phys. Lett . 51_, 221 (1977).

[18] Ravishankara, A. R. , Wagner, S., Fischer, S., Smith, G. , Schiff, R., Watson, R. T., Test,

G., and Davis, D. D., Int. J . Chem. Kinet .

,

in press (1978).

[19] Morris, E. D., Jr. and Niki, H., J. Phys . Chem . 75, 3640 (1971 )

.

[20] Atkinson, R., Perry, R. A., and Pitts, J. N., Jr., J. Chem. Phys . 67, 3170 (1977).

[21] Atkinson, R., Perry, R. A., and Pitts, J. H.

,

Jr., J. Chem. Phys . 38, 607 (1976).

[22] Atkinson, R. and Pitts, J. N . , Jr., J. Chem . Phys. 68, 2992 (1978).

[23] Atkinson, R., Darnall, K. R., and Pitts, J. N., Jr., J . Phys . Chem. , submitted for publication

(1978).

[24] Darnall, K. R., Winer, A. M., Lloyd, A. C.

,

and Pitts, J. N . , Jr., Chem. Phys. Lett . 44, 415 (1976).

[25] Perry, R. A., Ph.D. Thesis, University of California, Riverside, August 1977.

[26] Greiner, N. R., J. Chem. Phys . 5J_, 5049 (1969).

[27] Palmer, H. B., J. Chem. Phys . 64, 2699 (1976).

[28] Perry, R. A. Atkinson, R. , and Pitts, J. N.

,

Jr., J. Chem. Phys . 64, 5314 (1976).

[29] Perry, R. A., Atkinson, R., and Pitts, J.N.,

Jr., J. Chem. Phys . 67^, 5577 (1977).

[30] Simonaitis, R. and Heicklen, J., Int. J. Chem . Kinet . 5, 231 (1973).

'

[31] Howard, C. J., J. Chem. Phys . 65, 4771 (1976).

Julian Heicklen, Department of Chemistry, The Pennsylvania State University, University Park,

Pennsylvania 16802

9 All workers agree that at room temperature and atmospheric pressure the predominant reaction of

HO with C 2 Hit is addition. The only evidence for abstraction is given by Meagher and Heicklen (J_. Phys. Chem . 80, 1645, 1976) who photolyzed H 2 0 2 at I, 2537 & to produce HO radicals. They found C 2 H 5 0H

j, as the major product at high pressures and that

this product became less important as the pressure;

was reduced, as expected for the pressure-sensitive! addition reaction of HO with C 2 ^k- However they also found that CH 2 0 and HC00H were produced in a pressure-insensitive reaction. From

this they concluded that abstraction occurred 26 percent of the time for [C 2 HiJ y_ 2 to 5 Torr,

[H 2 0 2 ] 2l 2 Torr, and N 2 = 40 Torr. They further

| estimated that in the high pressure limit, this , fraction would further drop below 22 percent. Since then, the high pressure limit rate constant

j has been evaluated, and this fraction becomes 7 | percent at the high pressure limit.

This value is still higher than indicated by some other studies, and may be due to energetic

HO radicals in Meagher and Heicklen's system, since the photolyzing radiation provices an excess

of 23 kcal/mol over that needed to photo- dissociate H 2 0 2 . It is apparent that if the HO radicals become significantly energetic (either

thermally or by other means), and the pressure is low enough, then the abstraction reaction must

become dominant. Thus pressure and temperature

studies should be done to determine the conditions

when the two competitive paths are important.

William P. L. Carter, Statewide Air Pollution Research Center, University of California,

Riverside, California 92521

At the present time, the major uncertainties we|

; have found in developing and validating the mechanisms for the OH-olefin system concerns the rate constant for the decomposition of 6-substitut

ed alkoxy radicals, relative to the rate of their

reaction with 0 2 . For example, in the OH-propene

system, the question concerns the rates of

OH 0

I I CH 3 CH-CH2

6 OH

I I CH 3 CH-CH 2

CH3CHOH "+ HCH0

CH3CHO + CH 2 0H

(1)

(2)

relative to the competing reactions with atmo- spheric 0 2 , forming 3-hydroxy carbonyl products

(see Niki's figure 2

1 ). There appears to be some 1

Editor's note. Figure 2 in Dr. Niki's review

paper is a revised version provided by Dr. Carter: The change does not significantly affect Dr. Niki 's conclusions.

20
Iconflicts concerning this. The study of Niki et

h>. [1] on olefin-HONO-NO systems indicates that

poth reactions are fast, while we found, in Modeling [2] the more recent U.C. Riverside smog

Chamber data [3] that the acetaldehyde and formaldehyde yields in propene-containing NO

x-air

systems are better fit by models which assume that reaction (1) is slow and reaction (2) is fast.

[(The assumption that reaction (2) is much faster

than reaction (1) is consistent with the

jltheoretical estimates of Baldwin et al . [4]).

jAssuming that both reactions are fast results in

j [the model overpredicting the most recently

'determined acetaldehyde and formaldehyde yields

by ^ 25 to 50 percent.

' Our results [2] could be reconciled with those

of Niki et al . [1] if there were systematic

calibration errors in the determination of the acetaldehyde and formaldehyde yields in the UCR

smog chamber runs. However, the yields of these

.products monitored in n-butane-NO -air UCR chamber

runs using the same techniques [3j agree well with pur n-butane model predictions [2]. Since in the

| n-butane system, the photooxidation mechanism is

(less uncertain; this tends to indicate that the

(reported yields of these products in the recent

UCR chamber experiments are probably not in error.

L. Batt, Chemistry Department, University of Aberdeen, Aberdeen, Scotland AB9 2UE

By analogy with our studies on the decomposition and other reactions of alkoxy radicals we are able to make some prediction about the two unimolecular steps

6 OH

I I CH 3 - C - CH 2

•> CH 3 CH0 + CH 2 0H

H

OH 6

I I CH - C - CH 2

-* HCHO + CH3C-OH

given by Dr. Niki in his figure 2. The predictions indicate that the decompositions compete with difficulty with their reaction with oxygen,

if at all. (See more detailed comments in session on free radical chemistry).

It should be noted that our smog chamber

modeling results are completely inconsistent with

j both reactions (1) and (2) being slow. Not only

jdo models assuming this consistently underpredict

lacetaldehyde and formaldehyde, they also overpredict overall reactivity because both acetol and 2-hydroxy propanal , which would be the major

products if reactions (1) and (2) were slow, are expected to react with OH to form methyl glyoxyl

OH

I CH3CHCHO + OH

OH

H 2 0 + CH3C-CHO

OH 0 CH3C-CHO + 0 2 + H02 + CH3-C-CHO

jihose rapid photolysis would contribute significanTy to radical initiation [2].

eferences

jl] Niki, H., Maker, P. D. , Savage, C. M. , and Breitenbach, L. P., J. Phys. Chem. 82, 135

(1978). —

2] Carter, W. P. L. , Lloyd, A. C. , Sprung, J. L., and Pitts, J. N., Jr., Computer modeling of smog chamber data: Progress in validation of detailed mechanisms for the photooxidation of propene and n-butane in photochemical smog,

Int. J. Chem. Kinetics 11 , 45 (1979).

.3] Darnall, K. R. , Winer, A. M. , and Pitts, Jr.,

J. N. , A smog chamber study of the propene-n- butane-NO

x systems, in preparation (1978).

4] Baldwin, A. C, Barker, J. R. , Golden, D. M. , and Hendry, D. G. , J. Phys. Chem . 81, 2483

(1977).

Marciefe. Dodge, Environmental Protection Agency, Research Triangle Park, North Carolina 27711

We recently developed a mechanism for the propylene-0 3 reaction to use in our modeling

studies. The mechanism we formulated is based on the results of the two most recently published

studies of olefin-0 3 reactions. These are Herron

and Huie's study of the ethylene-0 3 reaction and the Niki et al_. study of the reaction of cis-2- butene with 0 3 in the presence of HCHO. The results of both of these studies can be explained

in terms of the Criegee mechanism, which is given below for the propylene-0 3 reaction:

CH3CH = CH2 + 03 ?

CH3CH CH2

CH3CH

O ^ HCH

CH3

.

We used the mechanism developed by Herron and Huie to explain the fate of the HCHOO radical. Our treatment of the CH 3 CH00 radical is based on

an analysis of the product yields obtained by Niki

et al_. in their study. In our model, a fraction

of the "hot" CH3CHOO radical is assumed to be collisionally-stabilized at atmospheric pressure.

The rest of the bi radical undergoes rearrangement

to form a "hot" acid and ester. The acid and ester subsequently decompose to various free

radical species.

This mechanism was used to model data collected

in our laboratory on the ozonolysis of propylene.

Four experiments were conducted in Teflon bags in air at atmospheric pressure. An example of the type of fits obtained when we modeled these data

is shown in figure 1. The simulated propylene

and 0 3 decay curves are in good agreement with

21
EXPERIMENTAL _

TIME, minutes

Fig. 1 Experimental and simulated results

using the new mechanism for olefin- ozone reactions.

the experimental profiles. The fits obtained for

the other three experiments were equally as good.

Although the new mechanism adequately explains

the observed decay of propylene and 0 3 , the mechanism favored by many modelers in the past

does not fit the data. In the last few years, many modelsrs have used a mechanism based on the O'Neal and Blumstein treatment of olefin-O^ reactions.

In this mechanism, the primary ozonide, after ringopening, can undergo a number of rearrangements,

the most likely of which is a-hydrogen abstraction

to form unstable hydroperoxides. These peroxides

can then fragment to an aldehyde and two free

radicals:

i I CH3CH -CH2

CH3CHCH

11 0OOH

I CH3CCH2

CH3CH + OH + HC

0

II

. CH3C

0

II OH + HCH

When this mechanism was used to model the data, we obtained the type of fit shown in figure 2. The simulated rate of propylene disappearance is significantly faster than the observed rate of

O OZONE

PROPYLENE — SIMULATED

1 1 1 EXPERIMENTAL

TIME, minutes

Fig. 2. Experimental and simulated results

using mechanism based on the O'Neal- Blumstein treatment of olefin-ozine

reactions.

loss. Clearly, the data do not support this trea ment of propylene-03 chemistry.

Although the mechanism developed in this study

adequately explains the observed decay of propylene and 0 3 , the results should not be construed as definitive. Additional work is needed in order to fully elucidate the mechanism

of ozone-olefin reactions.

W. Tsang, Center for Thermodynamics and Molecula

Science, National Bureau of Standards, Washington1 D.C. 20234 Many modeling studies appear to have as their

goal the matching of some particular set of experimental results (smog chamber data).

Considering the non-existence of a proper data base for such an effort, it is not clear what suc^ fits demonstrate. Certainly with the available number of adjustable parameters it does not take

a very ingeneous investigator to fit the data. To the unwary it may well appear that the entire problem has been solved. We know that this is no

the case and it would be more worthwhile to highlight disagreements and inability to fit the date This will immediately highlight the important

questions that must be settled.

H. Edward O'Neal, Department of Chemistry, San Diego State University, San Diego, California

92115Concerning the 0' Nea 1 -B'l umstein mechanism, it should be noted that in the original formulation

the rate of reaction from the mol ozonide to the Criegee intermediate was estimated using (for

an analogy) the then available t-butoxy radical I' decomposition rate constant. This is now known

to have an A-factor about 100 times higher than V

that used in the estimate. The Criegee reaction pathway is therefore corresponding more important!

and relative to the competing intramolecular

H-abstraction pathways, is now expected to be th&ij dominant process under many reaction conditions.

Gary Z. Whitten, Systems Applications, Inc., San Rafael, California 94903

In an ozone-olefin reaction, it is the yield 1 free radicals that is important and not necessari

the kind of radicals. Our recent modeling work | indicates that for ethylene the yield of free ra icals should not exceed about 10 percent. We we very pleased to see the results of Herron and Hui which confirm that estimate. In the case of the,; prophlene reaction we were pleased to see Niki'st

recent results which indicated a 30 percent radi-i cal yield, and Dodge's estimate of about 38 percent. In our air models we are using a value of

30 to 35 percent. Anything greater leads to a marked decay of propylene.

22
Julian Heicklen, Department of Chemistry, The

(ennsyl vania State University, University Park,

ennsylvania 16802

The reaction of 0 3 with olefins proceeds by

first forming the primary ozonide which then ecomposes to give either the Criegee zwitterion

lus a carbonyl compound or free radicals, owever 0 3 does react with olefins to form a eversible -rr-complex (E. Sanhueza, I. C. Hisatsune,

And J. Heicklen, Chem. Rev. , 76, 801, 1976).

[here is no evidence that this species plays any

fole in olefins containing only carbon and hydrojen. However this may not be the case with chlor-

[ated ethylenes, where kinetic evidence with

"HCICHCI gives a reaction rate law second order in

both CHC1CHC1 and 0 3 at very low pressures. This

j/as interpreted as meaning that the reversible

[-complex was the active species. This view is upported by the fact that the primary ozonides

cyere not seen in the reactions of 0 3 with cis- and

Jrans-CHCICHCI , CH 2 C1 2 , and C 2 CU (I. C. Hisatsune,

I H. Kolopajilo, and J. Heicklen, J. Am. Chem.

foe. , 3704, 1977). The possible role of the

.--complex should be considered in 0 3 reactions

nvolving substituted olefins.

Recommendations

Reactions of Ozone and Hydroxy! Radicals with Olefins

Reactions of olefins with ozone and hydroxyl

adicals are of fundamental importance for the hemistry of photochemical smog and although subtantial progress is being made in this field,

uch further work remains to be done. This work

hould involve both determinations of rate

onstants and mechanistic studies. The latter

hould be based on detailed product analysis,

implemented by computer modeling.

It is convenient to discuss ozone-olefin and

b'H-olefin reactions separately.

1 . Reactions of Ozone with Olefins

. Mechanism of ozone-olefin reactions

While significant progress is being made in the nvestigation of the mechanism of ozone-olefin factions, it is clear that the mechanism is not

et fully understood. Its full understanding is

f crucial importance for an understanding of the

ihemistry of photochemical smog. Some of the

hemistry involved in the 0 3 -olefin interaction in

he gas phase is probably the same or is similar

1 the chemistry of the oxygenated free radicals

leacting with 0 2 , as for example the radicals

brmed by addition of OH to olefins in the presence

| 0 2 . Further progress in this very difficult

'iled will require therefore imaginative studies

It only of ozone-olefin reactions but also of the

eactions of 0 2 with the oxygen containing free

adicals produced in these systems. Photolysis of

organic acids and esters and generation of selected oxygenated free radicals by other means

are examples of the techniques which could be utilized for this purpose in future work.

Recommendations :

1) Selected ozone-olefin reactions, including

cycloolefins, should be investigated under atmo- spheric conditions over a wide range of experimental parameters and with time resolved analysis of the concentrations of the reacting species and as many products as possible.

2) Techniques should be developed to generate

and study the chemical behavior of the Criegee

intermediates in the gas phase.

3) New approaches should be explored for studying the reactions of oxygen containing free radicals under atmospheric conditions to obtain information needed to understand the mechanism of ozone-olefin reactions in the gas phase.

4) The formation of aerosols induced by ozone- olefin reactions should be studied, the key intermediates isolated, and the critical chemical reactions studied.

B. Rates of ozone-olefin reactions

The phenomenological "rate constants" of the reactions of 0 3 with a number of simple terminal olefins in the gas phase, show good mutual agree- ment and describe well the rates of consumption of these olefins. However, their exact relation to the "true" bimolecular rate constants will only be resolved when the mechanism of ozone-olefin

reactions in the gas phase becomes fully understood. The "rate constants" for internal olefins measured in different laboratories show greater

discrepancies. These discrepancies may be largely

due to the fact that the range of experimental conditions has not been sufficiently broad to establish potential trends in the values. A better understanding of the reaction mechanism will no doubt also help to resolve these discrepancies.

The difference between the data obtained at high 0 3 -olefin concentrations in the gas phase and

in non-polar solvents those obtained at low 03- olefin concentrations in the gas phase and in non- polar solvents those obtained at low 0 3 -olefin concentrations under conditions similar to those

in the polluted troposphere is puzzling. Further work with the object of resolving this discrepancy, while not of the highest priority, could help in the understanding of the mechanism of the 03- olefin reactions in the gas phase. Recommendations :

1) Measurements of the rates of 0 3 -olefin reactions in the gas phase should be extended to cover a substantially broader range of experimental

condi tions

2) In view of the reportly very large 0 3 consuming effect of some olefins (e.g. cycloolefins, terpenes) in the atmosphere, their reaction rates should be redetermined.

23
2. Reactions of Hydroxyl Radicals with Olefins

A. Mechanism of OH-olefin reactions

The mechanism of the OH reactions with ethylene

and propylene in the absence of 0 2 appears now to be reasonably well understood. Hydroxyl radicals

add to these two olefins and there is little or no

H atom abstraction at room temperature. (A suggestion that there is approximately 8 percent H atom abstraction from ethylene at atmospheric pressure, possibly due to "hot" OH radicals, is given in a separate comment further below). The mechanism

of OH-olefin reactions in the presence of 0 2 , a process of crucial importance for the chemistry of photochemical smog, is unfortunately very in- completely understood.

Recommendations :

1) Study of the mechanism of the OH-olefin

reactions in the absence of 0 2 should be extended

to olefins other than C 2 Hi, and propylene,

especially to the olefins known to be present in the polluted atmosphere.

2) Very high priority should be attached to detailed studies of the OH-olefin reactions in the presence of 0 2 , especially for olefins known to be present in polluted atmosphere.

3) Studies of the OH-olefin reactions under atmospheric conditions in the presence of varying

amounts of NO (and possibly of other pollutants,

such as S02 , itc.) should be carried out with as detailed analysis of the reaction products as possi bl e.

B. Rates of OH-olefin reactions

Good experimental techniques for the determination of OH-olefin reaction rates are now available. However, caution has to be exercised to assure accurate determination of the very small reactant

concentrations used in some experiments and to establish the extent of the interfering secondary

reactions, in particular of the OH-free radical secondary reactions.

I 2) A set of accurate values of the rate constants of the OH reactions with selected olefins (including ethylene under conditions

similar to those in the lower atmosphere would be

j desirable in order to establish whether the rates

' are affected by oxygen in the air. The case of ethylene is of special interest in this respect because of a possibility of interception (and consumption by reaction) of the "hot" CH 2 CH 2 0H radicals by 0 2 . Such an interception could result

in an appreciable increase in the rate constant

of the 0H-C 2 H^ reaction in air relative to the value obtained in laboratory measurements in the absence of 0 2 .

3) An ongoing critical review of the rate constants would be very useful.

Rate constants for the simple olefins are now probably known to within ± 20 percent. The value 1 of the rate constant for C 2 H lt at 1 atm is probably also accurate within ± 20 percent. Rate constants for higher olefins and cycloolefins are less satisfactory, especially the values obtained I by the competitive technique. No values are available for some important naturally occurring

olefins such as terpenus and isoprene, although

the latter could be roughly estimated from the value of the rate constant for 1 ,3-butadiene. The range of the literature values of the rate constan

for acetylene is large (a factor of about 5-6) and further determinations are required.

Recommendations :

1) Further determinations of the rate constant

j

are required for higher olefins, cycloolefins,

isoprene, terpenes and acetylene.

!

24
Session II

Aldehydes
i.

r
TROPOSPHERIC CHEMISTRY OF ALDEHYDES

Alan C. Lloyd

Environmental Research and Technology, Inc. 2030 Alameda Padre Serra Santa Barbara, California 93103

This paper presents a survey of the current published literature on aldehydes, and to

a lesser extent, the other oxygenated hydrocarbons, as related to their role in modeling

the troposphere. Sources, ambient levels, photochemistry, and free radicals, reactions

of these substances are treated. Keywords: Aldehyde; free radical; photolysis; reactions; review; troposphere

1 . Introduction

Aldehydes are major products in the oxidation

)f hydrocarbons and play a rather unique role in

the photochemistry of the polluted troposphere.

: or example, they can contribute to photochemical

£mog, eye irritation, and odor problems. Their

iimportance has been recognized for over a decade

;(Leighton, 1961; Altshuller and Cohen, 1963;

Altshuller and Bufalini, 1965). While significant

iDrogress has been made in defining the photo-

:hemistry, kinetics, and mechanism of aldehyde

jhotooxidation , much remains to be learned about

their ambient concentrations as a function of

time, season and location. Since aldehydes,

Doth aliphatic and aromatic, occur as primary and

secondary pollutants and are direct precursors of

free radicals in the atmosphere, aldehyde chemistry

represents an important subject area. The

^understanding of this topic is necessary to meet

the objective of modeling tropospheric chemical

reactions. In this context, the major objective

of this paper is to consider the historical

interest in aldehydes; their sources and

atmospheric concentrations; the photochemistry,

kinetics and mechanism of their reactions and

finally to delineate current measurement needs

and recommend research priorities based on

(assessment of the current status of knowledge of

the chemistry of aldehydes in the troposphere.

In addition, the role of other oxygenated

hydrocarbons in tropospheric chemistry will be addressed briefly. Although aldehydes are the

main oxygenated hydrocarbons generally considered,

and will receive major considerations here, other

classes of oxygenated hydrocarbons merit

consideration and should be assessed in terms of

their involvement in the chemistry of the

|pol luted troposphere. Thus ketones, esters,

ethers and alcohols will be briefly considered

to assess their possible importance in modeling

the troposphere. The major areas of uncertainty

will be discussed and research priorities

jsuggested.

This paper is an attempt to survey the current

published literature on aldehydes (and, to a lesser extent, other oxygenated hydrocarbons) as the work relates to modeling the troposphere. It

is hoped that the discussion periods will extend

the coverage to include unpublished work, preliminary results, and peripheral studies which have

a direct bearing on the overall thrust of this

paper.

2. Previous Work and Importance of Aldehydes

Initial impetus for the interest in the role of aldehydes in photochemical air pollution

stemmed largely from the possibility that they were connected with eye irritation which became

a major phenomenon and problem in the Los Angeles

basin during the 1940's. However, an early Stanford Research Institute study (SRI 1950) concluded that "concentrations of aldehydes have rarely exceeded 0.2 parts per million by weight

and the high concentrations did not coincide with periods of eye irritation. This lack of correlation tends to indicate that aldehydes alone are not responsible for eye irritation." Subsequent

work indicated that acrolein was present on highly

polluted days and this compound is known to be a potent eye irritant (Los Angeles Air Pollution

Control District, 1950; Altshuller and McPherson,

1963; Scott Research Labs, 1969). Acrolein and formaldehyde were shown to be produced upon irradiation of dilute automobile exhaust and olefin-N0x mixtures (Schuck, 1957; Schuck and Doyle, 1959).

Aside from the possible relationship of aldehydes to eye irritation, it was subsequently

proposed (Leighton and Perkins, 1956; Leighton,

1961) that aldehydes could act as precursors to radicals which could either directly form oxidant

or oxidize NO to N0 2 . This possibility received

support from the results of several experimental

studies focused on the photooxidation of aldehydes

under laboratory and simulated atmospheric

conditions and generally employed formaldehyde

and the lower molecular weight aliphatic aldehydes

(Haagen-Smit and Fox, 1956; Altshuller and Cohen,

1963; Altshuller, Cohen et al . , 1966; Johnston

27
and Heicklen, 1964; Altshuller, Cohen et al .

,

1967; Cohen, Purcell et al . , 1967; Purcell and Cohen, 1967; Bufalini and Brubaker, 1969).

Recently Dimitriades et al . , (1972) and Pitts

et al . , (1976) carried out experiments in a smog chamber illustrating the effect of initial aldehyde concentrations on oxidant production

under simulated atmospheric conditions. Figure 1 shows the significant impact of initial aldehyde concentrations on ozone formation in a nine-hour

irradiation of a surrogate hydrocarbon mixture

(Pitts et al . , 1976). Thus an approximately 100 percent increase in initial formaldehyde concentration from 91 to 185 ppb increases the maximum ozone concentration by approximately

25 percent from about 0.39 to nearly 0.5 ppm in nine hours. Clearly, the rate of formation of

0 3 is enhanced but it is possible that the 0 3 maximum value would not be significantly increased

if the irradiations were carried out sufficiently

long.

2 4 6 8 TIME (hours) Fig. 1. Effect of added HCHO on ozone formation in long-term irradiations of surrogate mixture (from Pitts et al., 1976). Aldehydes can provide significant sources of radicals such as H0 2 , OH and RO2 which can influ- ence the rate at which photochemical oxidants are formed under ambient conditions. With the advent

of appropriate computer calculation facilities

to handle complex kinetic mechanisms, a number of workers demonstrated this effect by carrying out computer simulations of atmospheric chemistry

both with and without initial aldehydes (Niki,

Daby and Weinstock, 1972; Calvert et al . , 1972; Demerjian, Kerr and Calvert, 1974; Dodge and

Hecht, 1975; Levy, 1974; Whitten and Dodge, 1976; Graedel , 1976; Carter et al., 1978). Many of these calculations have focused on formaldehyde which photodissociates to produce significant

amounts of H0 2 radicals under ambient conditions.

Thus Demerjian et al . , (1974) have shown that

this route is the most important source of H0 2 radicals in the atmosphere.

Although there is some uncertainty attached to the quantum yields for photodissociation into radicals of HCHO as a function of wavelength

(vide infra), aldehydes are well established as important ingredients in photochemical smog formation.

The role of aldehydes as eye irritants and

radical precursors has been given above. An additional role for aldehydes is as precursors to the formation of peroxyacyl nitrates. These can

be formed by the reaction mechanism

RCHO + OH -> RCO + H 2 0

RCO + N0 2 * RCO3NO2

peroxyacyl nitrate

Peroxyacyl nitrate type compounds have been found in many parts of the world e.g., Penkett

et al . (1975) in England, van Ham and Nieboer

(1972) in Netherlands, in Japan (Akimoto and

Kondo, 1975) and in the U.S.A. (Stephens, 1969; Lonneman et al . (1976) )

.

3. Sources and Ambient Concentrations

Sources . There are primary and secondary

sources of aldehydes in the atmosphere. The primary sources are related to combustion and result from incomplete combustion in, for example,

internal combustion engines, diesel engines and stationary sources, such as incinerators, etc. (Altshuller et al . , 1961; Linnell and Scott, 1962; Elliot et al . , 1955). Automobiles are a significant source of aldehydes and the latter account for

up to one-tenth of the hydrocarbon emissions (Blacki 1977). Oberdorfer (1964) and Seizinger and Dimitriades (1972) have analyzed the individual I aldehydes emanating from pre-control led automobiles

Table 1 shows the percentage of aldehydes from automobile exhaust as determined by several workers (Oberdorfer, 1964; Fracchio et al . , 1967; ! Wodkowski and Weaver, 1970; Wigg et al . , 1972).

It is evident from these emission sources that formaldehyde is the largest aldehyde component. Similar but more extensive results are shown in table 2 which were obtained by Seizinger and Dimitriades (1972).

It can be seen that in addition to the saturated)

aliphatic aldehydes, acrolein--a potent eye irritant--is also present. In addition, benz- aldehyde and formaldehyde are produced, along with I alcohols, ethers and ketones. One would of course expect variations in the relative amounts of these compounds depending on the fuel used, e.g., see table 1

.

With the advent of hydrocarbon control measures j for automobiles, the aldehyde concentrations have been reduced along with the hydrocarbons. However, different control techniques apparently have varying effects upon the percentage reduction of aldehydes compared with the remaining hydrocarbons.;

Thus, Black (1977) shows interesting data for emissions from automobiles using thermal reactors, ! lean burn technology and catalysts of various

kinds to reduce hydrocarbons. Table 3 shows a comparison of absolute and relative hydrocarbon

class reductions for various automobiles employing

different hydrocarbon control systems. It is apparent that cars using the catalyst system, rather than the lean burn system, effect greater

reductions of aldehydes.

28
Table 1. Exhaust aldehyde analyses (adapted from National Academy of Sciences, 1976).

Fraction of total exhaust aldehydes, volume percent

Al dehyde Wi gg et al

(1972)

a Oberdorfer (1964) Wodkowski and Weaver (1970)

a Fracchio et al

Formaldehyde 66.7 72.5 70.2 59 9 69 3 72 9 Acetal dehyde 9.3 8.7 7.2 14 3 7 5 8 5

h Propionaldehyde 15.7 — 0.4

{' 0 0 7

( 6 4 Acrolein

(3.2 9.8 2 6 Butyraldehydes 4.3 0.4 3 0 1 0 1 7 Crotonal dehyde — -- 0.4 1 o 4 0 4

Val eral dehydes 0 4 Benzaldehyde 3.2 7.0 8.5 3 3 5 4 4 3 Tolual dehydes 1.9 7.2 5 9 3 1 Other 0.3 2.5 5 2 10 0 5 8^

TOTAL 100 100 100 100 100 100

a Exhausts from two different gasolines.

''Also includes acetone of unknown proportion.

Table 2. Oxygenates in exhaust from simple hydro- carbon fuels (from Seizinger and Dimitriades, 1972.). Concentration

Oxygenate range, ppma

Aldehydes are also emitted from some stationary

sources (EPA-AP42). These include power plants

burning oil and coal and sources such as incinerators, animal rendering facilities, or gasoline

and diesel engines operated at stationary source

facilities. Typical levels from stationary

combustion sources are given in table 4 (National

Academy of Sciences, 1976).

Secondary . Sources of aldehydes include the oxidation of hydrocarbons in the presence of NO under ambient atmospheric conditions. Figure 2 shows a concentration time plot for the formation

of secondary pollutants including aldehydes from

an irradiated hydrocarbon-NOx mixture under simulated atmospheric conditions. Major sources

in such systems are the reactions of ozone and OH radicals with olefins, and radical decomposition

products , e.g.,

0 3 + C 3 H 6

-y HCH0 + CH 3 0 + radical products

RCH0H + 0 Z

-* RCH0 + H0 2

In addition, aromatic aldehydes can be formed by the reaction of OH with aromatics, i.e., benz- aldehyde.

Ketones and hydroxyal dehydes or hydroxy- ketones

can obviously also be formed in such oxidation

systems

Ambient Concentrations . The major problem in measuring atmospheric concentrations of aldehydes

is the lack of an appropriate monitoring technique

applicable to this low concentration regime. For example, wet chemical techniques, such as the chromatropic acid method for measuring for- maldehyde, are subject to interferences and uncertainties in the results can be large.

Consequently, results of studies in the Los Angeles

area in the 1940's and 50' s may be subject to large error. For example, levels of 1.87 ppm in 1956 reported for total aldehydes by the Los Angeles County Air Pollution Control District are suspiciously high. More typical values of around

0.4 ppm were reported for lower molecular weight

aldehydes on days of significant air pollution

in Los Angeles in the late 1940's and early 50's

(Katz, 1961).

Acetal dehyde

b Propionaldehyde (+ acetone) Acrolein

Crotonal dehyde (+ toluene)

Tig! aldehyde

Benzaldehyde

Tolual dehyde

Ethyl benzaldehyde

o-Hydroxybenzal dehyde (+ do

aromatic)^

^ Acetone (+ propionaldehyde)

Methyl ethyl ketone

Methyl vinyl ketone (+ benzene)

Methyl propyl (or isopropyl) ketone

3- Methyl -3-buten-2-one

4- Methyl -3-penten-2-one

Acetophenone

Methanol

Ethanol x

C 3 alcohol (+ C 3 aromatic)

2-Buten-l-ol (+ C 3 H 8 0) Benzylalcohol

Phenol + cresol (s) 2,2,4,4-Tetramethyltetrahydrofuran Benzofuran

Methyl phenyl ether

Methyl formate Nitromethane CH3O

C5H5O C5H10O

0.8- 4. 9 2.3- 14. 0 0.2- 5. 3 0.1- 7. 0 <0.1- 0. 7 <0.1- 13. 5 <0.1- 2. 6 <0.1- 0. 2 <0.1- 3. 5

2.3- 14. 0 <0.1- 1. 0 0.1- 42. 6 <0.1- 0. 8

<0.1- 0 8 <0.1- 1 5 <0.1- 0 4 0.1- 0 6 <0.1- 0 6 <0.1- 1 1 <0.1- 3 6 <0.1- 0 6 <0.1- 6 7 <0.1- 6 4

<0.1- 2 8

<0. 1 <0.1- 0 .7 <0.8- 5 .0 <0.1 <0.1- 0 .2 <0.1- 0 .3

^from all test fuels.

Data represent unresolved mixture of propionalydehyde + acetone. Chromatographic peak shape suggests acetone to be the predominant component. ^Tuolene is the predominant component.

The C10 aromatic hydrocarbon is the predominant

component. ^Benzene is the predominant component.

The aromatic hydrocarbon is the predominant component.

29
Table 3. Automobile exhaust hydrocarbon (and aldehyde) emission patterns (from Black, 1977).

Total exhaust Percentage of total hydrocarbon, wt.%

Control Hydrocarbon, Total Vehicle system g/tni

.

Methane Paraffin Acetyl ene Benzene Aromatic Ethylene Olefin Formaldehyde aldehydes

1972

350 CID

non- catalyst

1.15 9.29 40.68 11.04 2.43 22.30 12.97 26.00 3.28 7.00

1974 Mazda

rotary

40 CID

thermal reactor

1.47 8.99 41.62 6.20 2.04 18.04 17.45 33.70 6.83 11.00

1975

Chev. Impala,

350 CID

single oxidation

catalyst

.25 15.44 57.09 2.72 1 .44 19.68 8.7.6 20.44 1 .60 3.64

Honda

CVCC

(proto.

)

91 CID

stratified

charge

(lean)

.28 4.64 37.82 7.29 1.79 22.24 15.11 32.46 6.00 11.29

Volvo

(proto)

130

three-way

reduction- oxidation catalyst

.16 19.91 56.59 4.73 1.47 21 .60 5.17 17.10 .94 3.25

Table 4. Typical emission of several classes of compounds including aldehydes from sta- tionary combustion sources (from Na- tional Academy of Sciences, 1976).

Emission, lb/ton of fuel Compounds Coal Oil Gas

Hydrocarbons 0.3 1.0 1.0 Aldehydes unknown 0.5 0.5

Formal dehyde 0.003 0.006 0.00? Organic acids 10 ^5 2

0.54

- 0.36

0.27

Fig. 2.

120 180 240 ELAPSED TIME (min)

Results of a typical smog chamber experi- ment (SAPRC evacuable chamber). Irradiation of a propyl ene-N0-N0 2 mixture in air. Initial experimental conditions--0.5 ppm propylene, 0.45 ppm NO, and 0.05 ppm N0 2 in 760 Torr of highly purified air.

Aldehydes have been measured in various parts

of the world, but the most extensive body of data exists for Los Angeles (Stanford Research Institute

1950; Renzetti and Bryan, 1961; Altshuller and McPherson, 1963; Scott Research Laboratories,

1969; California Air Resources Board, 1972). The studies of Altshuller and McPherson showed typical I formaldehyde and acrolein levels of ^ 0.04 ppm

and ^ 0.07 ppm respectively during September to November 1961. More recently, the California Air I Resources Board (1972) measured formaldehyde at levels up to 0.1 ppm, while daily average levels were around .035 ppm. Acetal dehyde exhibited an average concentration of .02 ppm, while no other j aldehydes or ketones were detected above their threshold of 0.1 ppm. The Air Resources Board found that aldehyde levels in the eastern part of the Los Angeles basin were significantly lower

than those in downtown Los Angeles: specifically, ; formaldehyde was found to average less than 0.2 ppm and acetaldehyde less than .015 ppm at Azusa. Figure 3 shows hourly concentrations measured at two locations in the Los Angeles area in 1968

by Scott Research Laboratories. Both locations

show sharp decreases in afternoon levels of total aliphatic aldehydes.

The advent of Fourier transform infrared spectroscopy (FTIR) has added a significant new dimension to the measurement of trace pollutants,

including aldehydes, in ambient air. The technique

is specific and sensitive. Hanst and coworkers

(1975) first applied the method in Pasadena in 1972. They measured HC00H at surprisingly high levels but detected no significant amounts of

HCH0. More recently, low levels of formaldehyde

(< 10 ppb) have been observed in the eastern

part of the Los Angeles basin by Tuazon et al. (1978a) also using FT-IR spectroscopy. These workers detected levels of formic acid up to 10 ppb and there was no obvious correlation between

the HCH0 and HC00H ambient concentrations.

However, this study has been largely superceded

by more recent results using improved absorptivi30
0.18 0.16

0.14

0.12 0.10

0.08 0.06

0.04

0.020

0.14 0.12

0.10

0.08

0.06

0.04

0.02

I 1 I I HUNTINGTON PARK / ,A\

ALIPHATIC ALDEHYDES FORMALDEHYDE . ACROLEIN

0.018 0.016

0.014

0.012

0.010

0.008 0.006

0.004 0.002 0

0.014

0.012 0.010

0.008

0.006

0.004

0.002

LOCAL TIME

Fig. 3. Hourly aldehyde concentrations at two Los Angeles sites, October 22, 1968 (from Air Fig. 5. Quality Criteria Document for Hydrocarbons,

1970).

HOUR OF DAY

Diurnal concentrations of formaldehyde at Newark, New Jersey for different days of the week, from June 1 to August 31 for the years 1972, 1973, and 1974 (from Cleveland

et al., 1977).

1400 1600 TIME (hours)

Fig. 4. Ambient concentrations of HCHO and HCOOH

as a function of time measured in Riverside, October 14, 1977 (from Tuason et al.,

1978).

ties for several species including HCHO and HCOOH

(Tuazon et al . , 1978b). Some of the results from

this 1977 study are shown in figure 4. The ambient levels of HCHO for this October 14, 1977

day (% 36 ppb) are significantly higher than measured in the previous study. This is partly

due to the improved absorpti vities used in the

1977 study. There is some evidence that the new

absorpti vities would also reduce the levels of

HCOOH reported by Hanst el al . (1973) in the Pasadena study (Winer, 1978).

In the last few years continuous measurements

for formaldehyde have been undertaken in certain

areas of New Jersey (Cleveland et al . , 1977).

This continuous monitoring showed a correlation

with vehicle traffic and a seasonal variation

with higher levels in summer than in winter. Peak formaldehyde concentrations were in the range of

14 to 20 ppb at four sites monitored. For example,

figure 5 shows formaldehyde levels reported for

Newark as a function of the day of the week. In Japan, Katou, (1972), observed high levels of the unsaturated aldehyde, acrolein. The average concentration measured was 7.2 ppb, but maximum

levels as high as 273 ppb were reported.

With the advent of FT-IR spectroscopy employed

by several groups of workers (e.g., Calvert et

al . ; Hanst et al.; Niki et al. and Pitts, Winer

et al . more reliable data for aldehydes should

become available for both ambient and simulated atmospheric conditions. Of necessity, the geographical area covered will be limited in the near future by the complexity and expense of the instrumentation.

4. Kinetics and Mechanism

This section is divided into two parts -- the first discusses the primary attack of radicals on aldehydes and the second part discusses the fate

in the atmosphere of the radicals produced.

The aldehydic hydrogen in aldehydes_is relatively

weak (C-H bond strength is 86 kcal mol l

, TrotmanDickinson and Kerr, 1975). Consequently, this hydrogen atom will be susceptible to attack by radical species present under atmospheric conditions. Possible species are 0(

3 P), 0(

1 D), OH, H0 2 , N0 3 and HSO^. Of these OH is likely to be by far the most dominant.

A. Radical and Atom Reactions with Aldehydes

OH Radical Reactions . OH is generally thought

to abstract an H atom from aldehydes -- chiefly

the aldehydric H atoms, i.e., the reaction

OH + RCHO * H 2 0 + RCO (1)

31
Niki and coworkers have carried out rate studies

for the largest number of aldehydes ranging from

the C1-C3 aliphatic aldehydes to benzaldehyde

(Morris and Niki, 1971; Morris et al., 1971; Niki et al . , 1978). The two former studies were carried

out at low pressure using a discharge flow-mass spectrometer technique for the generation of reactants and analysis of products respectively.

In their latest study, Niki et al . used the photolysis of HONO to generate OH radicals

in the presence of the aldehyde and C 2 H h or C 2 Di, at close to atmospheric pressure and monitored

the decay of reactants by Fourier transform infrared spectroscopy. In this way, rates of reaction

of aldehydes relative to ethylene were determined. These values were placed on an absolute basis using the appropriate rate constant for the OH reaction with C 2 H 4 at atmospheric pressure (Niki

et al . , 1978). These values are shown in table 5 with the modification that the rate constant for

the OH + C 2 Hi, reaction obtained by Atkinson et

al . , (1977) was used to reduce the relative values of Niki et al. (1978).

Also included are the recent results for HCHO

and CH3CHO obtained by Atkinson and Pitts (1978) using a technique yielding absolute rate constants. These latter workers used a flash photolysisresonance fluorescence technqiue and carried out

the first study of aldehydes over a range of temperature (299-426 K). Arrhenius activation energies for the two aldehydes studied are small with acetaldehyde exhibiting a negative value.

Table 5 allows a comparison among the results obtained by the various workers employing three different techniques. The agreement between the earlier work of Niki and coworkers (1971) and the most recent study of Niki et al. (1978) is excellent for HCHO and CH 3 CH0 but only fair for C2H5CHO. These results agree well with those of

Atkinson and Pitts (1978) for CH3CHO but are about

50 percent higher for HCHO.

If one assumes an atmospheric OH concentrations

of 10 6 radicals cm" 3

, the rates of decay of HCHO

and CH3CHO by reaction with OH are ^ 4.2 percent

and 5.8 percent h

1 respectively.

0 Atom Reactions . Attention here will be focused on ground state atomic oxygen, (0

3 P) since this more abundant than 0{ l D) in the lower

troposphere. 0(

3 P) reacts with aldehydes in the

same manner as OH, by abstracting the aldehydic

H atom. However, the reaction is a chain branching

one compared with a chain transfer reaction in the|j case of OH

0 + RCHO -> RCO + OH (2) Results of several studies of 0(

3 P) reacting with a variety of aldehydes are shown in table 6. No results are shown from purely high temperature

studies such as shock tubes.

The most extensive data are those obtained by

Singleton et al. (1977) for four aldehydes-acetalde

hyde, propionaldehyde, n- and iso-butyraldehyde. These workers used a phase shift technique and covered a temperature range of 298-472 K. They re4 ported that at the high end of their temperature

range, abstraction of the alky! group H atoms be- came significant for the higher molecular weight

aldehydes. However, under atmospheric conditions, abstraction of the aldehydric H atom is likely to

| dominate. The room temperature rate constants in ! table 6 show that there is generally good agreement

among the different workers for acetaldehyde but

1 significant differences for propionaldehyde and butyraldehyde. The technique used by Singleton et j al . , would suggest that their results should be morj reliable and should be used in any modeling studies'

Table 5. Rate constant data and Arrhenius parameters for the reaction of OH radicals with aldehydes.

10 12 x A Reactant cm

3 molec" 1 s"

ia

E cal mol"

10 12 x k cm

3 molec~'s~ at T K Reference

Temperature

range covered

HCHO

12.5 175 ± 300

CH 3 CH0

>6.6

14 ± 3.5 9.4 ± 1 .0 14.4 ± 0.8

15 ± 3.{

<20

6.87

C2H5CHO

C 6 H 5 CH0

-510 ± 300 16.0 ± 1.6

15.2 ± 1.6

31

20.8 ± 0.8

12.8 ± 0.8

300

298

299 298 ± 2

300

295 ± 2

299

298 ± 2

298

298 ± 2

298 ± 2

Herron and Penzhorn, 1969 Morris and Niki , 1971 Atkinson and Pitts, 1978

Niki et al . , 1978 (relative

to OH + C 2 H„ = 8.00 x 10" 12 (Atkinson et al .

,

1977))

Morris, Stedman and Niki, 1971

Cox et al., 1976 (relative to OH + HONO = 6.6 x 10" 12 (Cox, et al., 1976))

Atkinson and Pitts, 1978

Niki et al . , 1978 (relative to OH + C 2 H„ = 8.00 x 10 12 (Atkinson et al . , 1977))

Morris and Niki , 1971

Niki et al., 1978 (relative to OH + C 2 H„ = 8.00 x 10" 12 (Atkinson et al., 1977))

Niki et al, 1978 (relative to OH + C 2 H., = 8.00 x 10" 12 (Atkinson et al., 1977))

299-426 K

299-426 K

32 i
Table 6. Arrhenius parameters and rate constants for the reaction of oxygen atoms (0

3 P) with aldehydes.

Reactant

' A

(cm 3 molec" "s"

1 )

I (kcal mol (cm 3 k molec" 1 s" 1

) TK Technique Reference

HCHO 1

5 x 10 1 3 300 Discharge flow-mass ncl 1 Ull dilu rtril£MUril spectrometry (DF-MS) (1 9691

1 6 x 10" 1 3 300 DF-MS Niki, Daby, and Weinstock (1969)

1 5 x 10" 1 3 300 Discharge flow-product Mack and Thrush

analysis, epr and (1973)

chemi 1 umi nescence

pu run LH 3LHU 1 .20 x 10" 1

1 . 460 : t 0. 53 Hg sensitization—

.. ' . . jinQlstonGX dk l chemi luminescence (1977)

4 3 x 10 1 3 298

4 8 x 10" 1 3 298 Discharge flow-product Mack and Thrush

analysis, epr and (1974) chemi 1 umi nescence

2.3 x 10" 1

2 36 300-480 Discharge flow- Cadle and Powers

chemi 1 umi nescence (1967)

4 5 x 10" 1 3 298

5 0 x 10" 1 3 298 Hg sensitization- Cvetanovic (1956) product analysis

C 2 H 5 CH0 1 .30 x 10" 1

1.727 ± 0 066 298-472 Hg sensitization Singleton et al

chemi luminescence (1977)

7 0 x 10" 1 3 298

1 .41 x 10" 13 8 300-480 Discharge flow- Cadle et al. (1972) chemilumi nescence

2 3 x 10" 1 3 298

C3H7CHO 1 .66 x 10" 1 i 1 .702 ± 0 040 298-472 Hg sensitization- Singleton et al . (1977)

chemi luminescence

9 5 x 10" 1 3 298

2 5 x 10" 13 298 N0 2 photooxidation- Jaffee and Wan (1974) quantum loss of N0 2

i-C 3 H 7 CH0 1 .32 x 10" 11 .445 ± 0 091 298-472 Hg sensitization- Singleton et al . (1977) chemi 1 umi nescence

1

2 x 10" 12 298

LH2 - LHLHU 7.8 x 10" 12 0

-j r\f\ a 0 r\ jUU-4oU Discharge flow- Lad I e et a 1 . (

1 9/2

(acrolein) chemi luminescence

2. 7 x 10" 1 3 298

6.0 x 10" 1 L 2 84 one /too 296-423 Relative rate Gaffney, Atkinson, relative to 0 + and Pitts (1975) propylene

4. 9 x 10" 1 3 296

Croton3.3 x 10" 1 1 2 3 Discharge flow- Ladle et a 1 . ( ]9/4) dehyde chemi luminescence

8. 3 x 10" 1 3 298

6.4 x 10" 1 1 2 43 296-423 Relative rate Gaffney, Atkinson,

relative to 0 + and Pitts (1975)

1.0S x 10" 12 296 propylene

The three major studies of formaldehyde at room

temperature show excellent agreement. The rate constant appears to be about one third that of acetaldehyde which reflects the weaker aldehydric

H bond in acetaldehyde.

If one assumes an 0(

3 P) atom concentration of

10 5 atoms cm

-3

, for the lower troposphere, the

rates of reaction of HCHO and CH 3 CH0 with 0(

3 P) are 5.4 x 10" 3 and 1.5 x 10" 2 percent h" 1

respectively. Thus this route will be unimportant

for the atmospheric chemistry of aldehydes.

Measurements have been reported for the un- saturated aldehydes acrolein, CH 2 = CHCH0 and crotonal dehyde, CH 3 CH = CHCH0. Cadle et al . (1972,

1974) used a discharge flow technique over the approximate temperature range of 300-480 K. The values shown in table 6 are in reasonable agreement

with those of Gaffney et al. (1975). The latter

33
workers used a relative rate technique using the mercury photosensitized decomposition of N2O at 2537 A to generate 0(

3 P) atoms. Product analysis

was by gas chromatography.

H0 2 Radicals . No room temperature rate constants

have been reported for the reaction of H0 2 radicals

with aldehydes, although Baldwin and coworkers

(1972) have_obtained a value of 1.6 x 10" 15 cm

3 molecule"^" 1 at 773 K for formaldehyde. In addition, Hendry and Mabey (1973) reported a value for HCHO of 7.1 x 10" 18 cm

3 molecule 1 s" 1 at 373 K. In order to obtain an estimate for the rate of reaction of

H0 2 + HCHO -* H 2 0 2 + CHO (3)

one may employ the expression suggested by Lloyd

(1974) of 1.7 x 10" 12 exp(-4700/T) cm

3 _ molecule"^" 1 to produce k 3 = 2.8 x 10 18 cm

3 molecule'V 1 at 298 K.

Combining this value with a typical H0 2 concentration of 10 9 radicals cm

3 for the_polluted

troposphere one obtains a value of 1 x 10 3 percent h" 1 for the rate of disappearance of HCHO

by reaction with H0 2 . This rate would be about

an order of magnitude smaller in the unpolluted

troposphere.

Although measurements of the rate constant for

H0 2 reacting with HCHO would be desirable from a scientific viewpoint, unless current measurements

at higher temperatures are grossly in error, it does not appear from an atmospheric modeling view- point that this reaction plays a significant role

in the chemistry of the polluted troposphere.

Alkoxy Radicals . Kelly and Heicklen (1978) have recently measured the rate constant for the lowest molecular weight compound in the series,

methoxy radicals reacting with acetaldehyde

CH3O + CH3CHO + CH3OH + CH3CO (4)

This is a radical transfer reaction. The authors photolyzed azomethane in the presence of acetal- dehyde and oxygen at 298 K and from a product

analysis obtained U/k 5 = 1.4 ± 2.8 where k 5

CH3O + 0 2

-> CH 2 0 + H0 2 (5)

Using the value of 6 x 10" 16 cm

3 molecule" 1 s"

1 obtained by Barker et al . (1977) for k 5 , the authors guoted = (8.3 ± 1.7) x 10" 15 cm

3 molecule 1 s" 1 at 298 K.

The steady state concentration of CH 3 0 in the lower polluted troposphere is around 5 x 10 6 radicals cm" 3

. Thus the rate of disappearance of acetaldehyde by reaction with_CH 3 0 in thejower

troposphere is about 1.5 x 10" 2 percent h

1

.

N0 3 Radicals . Morris and Niki (1974) studied

the reaction

N0 3 + CH3CHO HNO3 + CH3CO (6)

using a long path length IR cell operated at 300 K and near atmospheric pressure. The results from studying the reaction of mixtures of N 2 0 5 and

CH 3 CH0 were interpreted using numerical integratio!

for the participating reactions. A value of

k 6 = 1.2 x 10' 15 cm

3 molecule" 1 s"

1 at 300 K

was obtained by varying k 6 until a good match of

j the calculated and observed N 2 0 5 decay was obtained.

_Assuming an N0 3 concentration of 10 9 radicals 1 cm

3 for the polluted troposphere, one can estimat

that the rate of disappearance of CH 3 CH0 by reaction 6 is 0.4 percent h

l

.

HSO., Radicals . This radical is given some consideration here although no experimental data

are available for the relevant reaction:

HS0* + RCH0 * H 2 S0„ + RC0 (7)

1

1

Benson (1978) has suggested that the HSCK radical I could be more reactive than R0 radicals in either I adding to the double bond of olefins or in H- kr abstraction reactions. Since aldehydes have a relatively weak C-H bond, reaction 7 is a logical

p candidate to consider in HC-N0 -SO photooxidation

systems

.

r HSOk can be formed in the sequence of reactions

OH + S0 2 + M * HSO3 + M

HS0 2 + 0 2

-> HSO5

HS0 5 + NO * HS0 4 + N0 2

Using AH

f (HS0-) = -125 kcal mol 1 (Benson,

1978) one cSn estimate that reaction 7 for HCHO

is 17.1 kcal mol" 1 exothermic. (The similar

reaction for OH radicals reacting instead of HSOt,

jj is 32.30 kcal_mol

_1 exothermic and for CH 3 0 is 17.5 kcal mol" 1 exothermic). Thus, by comparison with the sole experimental _measurement of k^, one \ would expect k 7 ^8.3 x 10~ 15 cm

3 molecule" 1 s

1 at 300 K. However, this reaction should be studied experimentally since thermochemistry is not always

a reliable guide for estimating kinetic data.

B. Atmospheric Reactions of Radicals Produced

from Attack of Radicals on Aldehydes

We have seen above that radicals of the form

j RC0 are produced from the reaction of atoms and free radicals with aldehydes. In this section,

the subsequent fate of these radicals under atmo-

' spheric conditions will be discussed. Differentia-1 tion is made between the acyl radicals and their aromatic equivalents since there is evidence

(Niki et al . , 1978) that the radical produced from benzaldehyde react differently from their aliphatic^

equivalents

.

The simplest acyl radical is formyl produced

from formaldehyde. Three reaction paths are possible for its reaction under atmospheric

conditions. These are:

HC0 + 0 2 > H0 2 + CO (8) HC0 + 0 2 + OH + C0 2 (9) HC0 + 0 2 + M * HCO3 + M (10)

\
The following results support the conclusion

that reaction (8) is the main route for this reaction

- the reaction obeys second order kinetics at low pressures (Washida et al . , 1974),

- the rate constant for HCO + 0 2 is independent

of pressure over the range 20-500 Torr

(Shibuya et al . , 1977),

- the direct identification of H0 2 formation

from HCO by laser magnetic resonance (Radford

et al . , 1974),

- the observation of H0 2 formation from HCO +

0 2 at 1 atmosphere (Hunziker, 1975).

However, results from recent studies by Osif

and Heicklen (1976) and Niki et al . , (1977) suggested that reaction (10) was the dominant

| pathway. Both of these studies used the CI atom

[sensitized decomposition of formaldehyde in the

I presence of 0 2 . They assumed that formic acid

| formation was a good indicator of reaction (10),

I, since HC0 3 radicals would be converted to HC0 2 radicals and subsequently HCOOH by H-abstraction

from the aldehyde, thus

HC0 3 + HCO3 > 2HC02 + 0 2 (11 )

HC0 2 + HCHO -> HC0 2 H + HCO. (12)

I Osif and Heicklen measured ki 0 /k 8 ^ 5 ± 1 independent of pressure over the range studied (70-700

Torr) while Niki 1 et al . , obtained a value > 9 for

the same ratio at atmospheric pressure. Hanst

and Gay (1977) also used the CI atom catalyzed

oxidation of HCHO. They irradiated low concentrations of CI 2/HCHO/NO2 mixtures in 1 atmosphere of

air and the analysis performed using FT-IR spectroscopy. From the small yield of HCOOH and the observations of peroxynitric acid H0 2 N0 2 , they concluded that reaction (10) was unimportant in their system. Horowitz et al. (1978) have subsequently

suggested that formic acid can be formed by routes

other than those involving reaction (10), specifically

OH + HCHO (H0CH 2 0) -* HC0 2 H + H (13)

j H0 2 + HCHO + (H0 2 CH 2 0) -> HC0 2 H + OH (14)

•they state that reactions (13) and (14) are 22 and

60 kcal mol 1 exothermic respectively. These

workers photolyzed mixtures of HCHO at 3130 K at

low pressures in the presence of 0 2 and added C0 2 ,

; and measured the quantum yields of formation of

H 2 , CO and C0 2 and the loss of 0 2 . A lower limit

for k 10 of 1.21 x 10~ 2B cm

6 molecule" 2 s"

1 was obtained from the measured ratio k 10 M/(ki 0 M + k 8 ) > 0.049 ± 0.017 (obtained from experiments using

Niki (1978) suggests that formation of HCOOH in his system may be explained by Reaction (14) as suggested by Horowitz et al . (1978). This would

reduce the importance of Reaction (10) in the system used by Niki et al . (1977).

p HCH0 = 8 Torr and p n

= 1-4 Torr) and Washida

et al.'s (1974) value 2 of k 8 . Contrary to Osid

and Heicklen (1976) and Niki et al. (1977), Horowitz et al . conclude that reactions (8) and

(10) assume about equal importance under atmo- spheric conditions. Clearly, further work is needed to clarify this discrepancy.

Table 7. Rate constants for HCO + 0 2 + H0 2 + CO at 298 K. Rate constant, k 8 Pressure Reference

(cm 3 mol ec"

1 s" 1 ) x 10 12 Torr

5 7 ± 1 2 4 Washida, et al. (1974)

6 0 ± 0 9 20 Shibuya, et al. (1977)

5 3 ± 0 7 530 Shibuya, et al. (1977)

3 8 ± 0 6

? More, quoted in Shibuya (1977)

Table 7 shows the generally good agreement for results of rate constant determinations of k 8 . Washida et al . (1974) used a photoionization mass spectrometer coupled to a flow system to obtain

k 8 = (5.7 ± 1.2) x 10" 12 cm

3 molecule" 1 s" 1 at room temperature. Recently, Shibuya et al (1977) generated HCO radicals in the absence and presence of 0 2 by the flash photolysis of CH3CHO

CH3CHO + hv(> 2000 A) •> CH 3 + HCO.

From an analysis of the behavior of HCO radical decay, a_value of k 8 = (5.6 ± 0.9) x 10" 12 cm

3 molecule" 1 s"

1 at 298 K was obtained. This value

is essentially independent of pressure (see table 7) and is in excellent agreement with that

of Washida et al . (1974) but larger than the value obtained by Moore quoted in Shibuya et al

Under atmospheric conditions, acyl radicals other than formyl are generally assumed to react with 0 2 by addition:

RC0 + 0 2 RC0-0 2 .

Subsequent reactions in the polluted atmosphere

with NO and N0 2 occur

RC0-0 2 + NO + N0 2 + RC0 2 - R + C0 2 (15)

RC0-0 2 + N0 2 * RC0-0 2

• N0 2 peroxyacyl nitrate

The most commonly studied member of the peroxyacyl nitrate family is peroxyacetyl nitrate (PAN) (Stephens, 1969; Pate, Atkinson and Pitts, 1976; Hendry and Kenley, 1977; Cox and Roffey, 1977). The kinetics and mechanism of PAN formation and

thermal decomposition have been discussed recently

(Pate et al . , 1976; Hendry and Kenley, 1977; Baldwin et al . , 1977 ; Carter et al . 1978) and is beyond the scope of this paper.

Niki et al . (1978) note that, based on OH reactivity, aliphatic and aromatic aldehydes,

35
should be equally efficient at converting NO to N0 2 . However, they note that smog chamber studies

show that benzaldehyde is far less reactive than

the aliphatic aldehydes in producing ozone (Dimitriades , 1974). They suggest that radicals

formed by H abstraction from benzaldehyde are efficient NO scavengers. Support for these suggestions would significantly aid the understanding and computer modeling of aromatic hydrocarbon

photooxidation.

5. Photochemistry of Aldehydes

The photodissociation of aldehydes is an import- ant radical generation mechanism in the formation

of photochemical air pollution (Leighton, 1961; Altshuller and Bufalini, 1965, 1971; Pitts, 1969; Calvert et al . , 1972; Demerjian et al . , 1974; Levy, 1974; Hecht, Seinfeld and Dodge, 1974; Dodge and Hecht, 1975; Whitten and Dodge, 1976).

2000 2200 2400 2600 2800 3000 3200 3400 3600 3800 WAVELENGTH, A

Fig. 6. Absorption spectra of formaldehyde (1), acetaldehyde (2), and propionaldehyde (3) (from Calvert and Pitts, 1966).

Figure 6 shows the absorption spectra for some common aldehydes which illustrates that they will absorb well beyond 3000 A. The two reactions of most significance can be generalized in terms of

a radical and molecular route:

RCHO + hv

R + HCO

RH + CO

The radical route is the more important one for modeling tropospheric chemistry. The rate constant for a particular primary process is an important quantity in assessing the importance of

the process in the atmosphere. It is given by

k.(0,h) =E J(X,G,h) • o.(X)

1 X

where 0 is the solar zenith angle, h is the height

above ground, X is the wavelength, J is the actinic

flux, a is the cross section and 0 is the quantum

yield for species i. Of these parameters, the quantum yield $ as a function of wavelength has been subject to major uncertainty; for example,

singificant differences among experimentally

determined values have been reported for formaldehyde (McQuigg and Calvert, 1969; Sperling and Toby, 1973).

Quantum Yields for HCHO photolysis as a Function of wavelength . Considerable attention

has been given to formaldehyde photolysis in recent years, partly because of its importance in photochemical air pollution. There appears to be general agreement that:

a. the major final products are H 2 and CO (in the absence of 0 2 )

,

b. the primary reaction paths are

HCHO + hv -y H + HCO

H 2 + CO

(17) (18)

However, differences have arisen in the determi nations of the variation of the ratio of the quantum yields of reactions (17) and (18), $i 7 /$i 8 as a function of wavelength. The earlier studies

(pre 1975) showed no consistent trend with wavelength variation (Gorin, 1939; Klein and Schoen,

1956; DeGraff and Calvert, 1967; McQuigg and Calvert, 1969; Sperling and Toby, 1973). The results of these studies have been superceded by results from more recent and definitive studies

and can be discounted (Horowitz and Calvert, 1978; Lewis and Lee, 1978). Consequently, only the recent studies (post 1975) will be discussed here and emphasis will be on the radical production

route, reaction (17).

As pointed out by Horowitz and Calvert (1978),

all the recent studies show that the quantum

yield for the radical process, $i 7 , increases more steeply with decreasing wavelength below about

3400 A. This is illustrated in figure 7. This figure is similar to that given by Horowitz and Calvert (1978b) but it has the additional results of Moortgat et al . (1978). These results give

slightly higher values of $ i7 and show a lesser tendency to level off below about 3200 A.

The results for $i 7 reported by Lewis et al

.

(1976) should be corrected by a factor of 1.89 according to Lewis and Lee (1978). Lewis et al. irradiated a mixture of HCHO and NO with mono- chromatic radiation (= 1 A bandwidth) from a tunable laser and measured the intensity of chemi luminescence produced from excited HNO*

formed in the three-body recombination H + NO + M.

H atoms were produced in reaction (17). Relative

values of $i 7 were obtained and were converted to absolute values using <h 7 = 0.36 ± 0.04 at 3035

& determined in a separate experiment (Lewis and Lee, 1976). It is this value which has been redetermined by Lewis and Lee (1978) and increased

to 0.68 ± 0.05. This change brings the earlier

results of Lewis et al . (1976) into line with the most recent determinations discussed below.

Horowtiz and Calvert (1978a and b) have recently determined $ 17 at 3130 A and over the wavelength

range 2890 to 3380 & at 398 K. They obtained $ 18 values by measuring the quantum yields for H 2 production, $^ 2

> in the photolysis of HCHO- isobutene mixtures assuming that the high concen- trations of isobutene scavenged the H atom production by reaction (17). Values of <h 7 and

$i8 were obtained from *H2 in the photolysis of pure HCHO. These authors found f 17 and $ 18 to be essentially 1 until the longest wavelength

studied (3380 A). $ J7 was found to increase from

0 at 3370 A to = 0.7 at 3175 A in general agreement

ic ;;i

D

h

36
with other recent studies. Results of experiments

at 3380 A with added oxygen lead the authors to conclude that little if any dissociation of HCHO

into radicals occurred at 3380 A and longer wavelengths, in contrast to earlier results.

Marling (1977) photolyzed 4 Torr of HDCO using either a high pressure mercury arc coupled with a monochromater of monochromatic laser jrradiation

in the wavelength range 3040 to 3530 A. He measured the relative yields of H 2 , D 2 and HD as

a function of wavelength using a mass spectrometer

and noted that the (H 2 + D 2 ) yield is a measure of

radical production via reaction (17), since H 2 and

D 2 can only be formed following reaction (17). Marling found that radical production reached

55 percent at wavelengths less than a 3200 A, but

no measurements of the absolute decomposition

yield were reported. Horowitz and Calvert (1978) have placed Marling's results on an absolute basis

using their value of $ i7 = 0.76 at 3030 A and applied the ratio of 0.55 measured at 3040 A by

Marling. They find that, within the uncertainties

introduced by isotopic differences in the HCHO and

HDCO molecules, Marling's reduced results agree reasonably well with their own (see fig. 7 and table 8).

Clark (1976) photolyzed HCHO in the presence of

NO using a tunable laser. The NO was employed as

a radical trap for HC0 radicals and H atoms and

j was used in both small and large quantities and

I the effect on $y 2 and $

c q measured. He observed

I the same sharp increase in $i 7 in the 3200 to 3300 & range as other recent studies, and concluded

that $17 + $i8 = 1.0. Different results were

obtained in experiments with low NO and high NO. Clark assigned more credibility to results obtained

with low NO and suggested that high NO enhanced

HCHO decomposition by reaction (17). However,

Horowitz and Calvert (1978) have questioned this interpretation as a consequence of the results of their studies with added NO which showed little

effect on $H2. They conclude that Clark's results

for $

H2 at high NO are more appropriate estimates

of $i8, and find that when these results for $i 8 are used, Clark's results for $ 17 are in excellent

agreement with their own studies. In view of the

results of other recent studies, the suggested

reworking of Clark's results appears valid.

Moortgat et al . (1978a) have studied the photolysis of HCHO at several wavelengths in the range

2700 to 3600 ft. They employed two types of experiment -- HCHO in a mixture of N 2 with a small amount of C 3 H 6 added, and HCHO in a 20:80 mixture

of 0 2 and N 2 -- and both systems were operated at atmospheric pressure. Consequently, results from

the latter should be directly applicable to modeling the lower troposphere. A Xenon arc mono- chromater was used to isolate the desired wavelengths. Measurements of the ratio of H 2 and CO production yielded values for the ratio

<J

> ie/$i7 + $is- Subsequent work (Moortgat et al.,

1978b) reported values for $i 7 + $ib, and consequently, $i 7 . Although initially the values

for $ 17 obtained from the C 3 H 6 added experiment

were lower than those for the N 2 :0 2 system,

subsequent allowance for a small contribution to

the H 2 formation by H abstraction from C 3 H 6 (Moortgat, 1978) gave excellent agreement between

the two sets of results as shown in table 8.

Examination of figure 7 shows that results of recent studies show a consistent trend although

there remains some scatter in the results of different studies. However, it is possible to draw a line or band which incorporates most of the results within their experimental error. This significant improvement in our knowledge of $i 7 is very helpful in improving our capability to model atmospheric chemistry in the polluted

troposphere.

o

*

LEWIS ET AL 1976 (CORRECTED)

LEWIS ET AL 1978 HOROWITZ AND CALVERT 1978 MARLING 1977 CLARK 1976 MOORTGAT ET AL 1978

3000 3100 WAVELENGTH, IA)

Fig. 7. Primary Quantum Yield, $ 17 , for process

HCHO + hv -+ H + HCO, as a function of wavelength. Data of Marling and Clark

plotted using interpretation of Horowitz

and Calvert.

In order to compare the rate of photolysis

with the depletion of formaldehyde by radical processes one can calculate a photolysis rate of ^ 13 percent h" 1 for a solar zenith angle of 20° and using the value of k i7 given by Horowitz and Calvert (1978b).

Photolysis of Acetaldehyde . Acetaldehyde is commonly used as a surrogate for aldehydes of higher molecular weight than formaldehyde. Its absorption spectrum is shown' in figure 6, which is taken from Calvert and Pitts (1966). As with formaldehyde, major uncertainty is concentrated

in the quantum yields of the primary processes

(Calvert and Pitts, 1966). Parmenter and Noyes (1963) carried out emission

studies and Archer et al . (1973) used the triplet

state induced cis-trans isomerization of butene-2

to study the primary processes in acetaldehyde

photolysis. These studies have been summarized

by Weaver et al . (1976, 77). In a comprehensive

study, these latter workers obtained results which

are consistent with the previous studies. Weaver

et al photolysed CH 3 CH0 vapor at 3130 # in the presence of 0 2 or 0 2 - N 2 mixtures at 298 K. The products formed at a function of pressure and added 0 2 were measured over the pressure range 20

to 640 Torr. Weaver et al . postulated the following

37
Table 8. Wavelength dependence for the quantum yield *, of HCHO photodecomposition into H and HCO.

Wavelength Quantum Yield $

'

Lewis Lewis Horowitz Marling Clark c Moortgat et al. et al. a and Lee and Calvert

2767 0.48 0.44 2754 0.68 ± 0.10

2840 0.51 2841 0.65 0.66

2882 0.64 0.64 ± 0.10

2890 0.701 2894 0.711 2930 0.740

2934 0.49

2950 0.80 0.81 2982 0.62

2991 0.70

(0.48) 3030 0.760

3035 0.68 0.68 ± 0.05

3039 0.84 0.81 3040 0.760

3050 0.49

3088 0.60

3130 0.735

3140 0.70 0.760

3163 0.80 0.80

3166 0.62 3172 0.64

(0.42) 3175 0.692 3180 0.636

3195 0.635

3210 0.51 0.554

3230 0.519

3250 0.460 0.456

3260 0.42 0.442

3264 0.48

(0.42) 3267 0.48 0.49 3270 0.438

3296 0.28

3298 0.38

(0.31) 3300 0.330 3310 0.097 3324 0.097

3335 0.212 3340 0.113

3360 0.048

3378 0.111 3380 0.00

3392 < .19 0.01 (-0.03) 3402 0.13 0.09

3550 0.04 0.03

jjoriqinal results multiplied by 1.89 as recommended by Lewis and Lee (1978). Results quoted by Horowitz and Calvert from Marlings experimental data.

c Values quoted are those from high NO data reinterpreted by Horowitz and Calvert. Numbers in parentheses are those given by Clark for low NO pressures.

reactions to account for their results

CH 3 CH0 + hv > 1 [CH 3 CH0]

n (19) (quenchable part of excited singlet

state) + ^CHsCHO]' (20) (non-quenchable part of excited

singlet state)

^CHaCHO]' - CH 3 + HCO (21)

-v CH3CHO (22)

1 [CH 3 CH0]

n + M - ^CHaCHOJo + M (23) (vibrationally equilibrated electronicallyexcited single state)

1 [CH 3 CH03o +

3 [CH 3 CH0] (24)

3 [CH 3 CH0] + 0 2 > [complex] (25)

[complex] - CH 3 + CO + H0 2 (26)

Weaver et al . found no CHi* produced at 3130

in the presence of NO and so concluded that the other primary process, reaction (27), was negligible at 3130 A.

CH 3 CH0 + hv ->- ChU + CO (27)

The quantum yields obtained are shown in table which is taken from Weaver et al. (1976, 77).

Reactions (24) through (26) which predict a pressure dependent reaction between the excited

38
Table 9. Quantum yields of the primary processes

in acetaldehyde photo-oxidation as a function of excitation wavelength

(Weaver et al ., 1976).

A (fl) 0{CQ + CH„) 0{CH 3 + HCO} gHtriplet}

3340 0 0 l-0

d

3130 0 0.05

b 0.84

e

2967 — — 0.59<[ 2804 0.1

a <0.30c 0.48

2654 0.28

a 0.36

a

2537 0.64

a 0.36

a

a From Calvert and Pitts (1966).

bThis work and Archer, et al . (1973). Calculated from the total quantum yield of the free radical process (0 {CH 3 + HCO} = 0.39, Calvert and Pitts (1966); the fraction of that from the triplet, 0.18; and 0{triplet} at 2804 \ = 0.48. From Parmenter and Noyes (1963).

e From Parmenter and Noyes (1963) and Archer, etal. (1973).

triplet state and 0 2 , form the novel part of Weaver et al.'s results. Although they were able

to fit their results with their suggested

mechanism, the quantum yields obtained are subject

to uncertainty since the mechanism for secondary

reactions may affect the results. For example, it was assumed that HCO reacted with 0 2 by addition

(contrary to the conclusion of our discussions

above) and the reaction of CH 3 0 2 with H0 2 had to

be omitted in order to fit the data.

Weaver et al calculate the rate of CH 3 CH0 photolysis in the atmosphere for a solar zenith

angle of zero (i.e. the maximum photolysis rate)

and find the rates for reactions (21) and (26) to

be 2.8 x 10" 6 and 8.7 x 10~ 6 s" 1 respectively.

They report the overall rate constant for all free

radical processes to be 2.3 x 10~ 5 s a 8.3 percent

h -1

. This may be compared with the role of depletion of CH 3 CH0 by reaction with_0H radicals

(assumed [OH] = 10 G radicals cm" 3

) of 5.8 percent

h" 1.

6. Other Oxygenates

In addition to aldehydes, there are several other classes of oxygenated hydrocarbons known or suspected to be present in urban areas of the troposphere. These include ketones, alcohols,

esters and ethers. Their possible role in atmospheric chemistry will be discussed briefly.

Discussion of radical reactions with these

oxygenates will be limited to the OH radical.

This was shown to be the major species attacking

: aldehydes and is likely to be the most important

I intermediate for the remaining oxygenates. Of

1 the four classes of oxygenated hydrocarbons

I mentioned above, only ketones undergo photolysis

I under ambient conditions and this process is discussed later.

Sources and Ambient Concentrations of Ketones, Alcohols, Esters and Ethers . All of these classes

of compounds are used in commercial solvents

(Burnelle et al . 1966; Wilson and Doyle, 1970; Levy and Miller 1970; Laity et al . 1973). These compounds occur in paints, degreasing solvents,

etc. Additionally, automobile exhaust contains

small quantities of all the above mentioned classes

as illustrated by table 2. Ketones and alcohols

are also formed via secondary reactions in the atmosphere in a manner similar to that for the

al dehydes

Data on ambient concentrations of these compounds are scarce and the only available quanti- tative measurements appear to be for selected

ketones. Acetone measurements of 0.3 to 0.9 ppb have been reported by Robinson et al. (1973) for remote areas of California, Idaho, Vermont and Washington. Methyl ethyle ketone (MEK) has been observed in Riverside, California at concentrations

of 1 to 6 ppb (Stephens and Burleson, 1976). Smoyer et al . (1971) report the detection of

94 ppm of MEK in ambient air near a chemical re- clamation plant in Maryland. Grob and Grob (1971) detected acetophenone in ambient air in Zurich, Switzerland.

Reactions of Oxygenated Hydrocarbons with the Hydroxy! Radicals . The results of studies of the reactions of OH radicals with individuals oxygenated hydrocarbons are summarized in table 10, which

is a modified version of that given by Atkinson

et al . (1978). Most of these determinations are

for one temperature (around 300 K) and were obtained

using a relative rate technique. Absolute values were obtained as indicated in table 10.

The rates of disappearance of these oxygenates

due to reaction with OH radicals are given in table 11, assuming an [OH] of 10 6 radicals cm" 3

. The rates given for the ketones will underestimate

their disappearance rates in the atmosphere since they can photodissociate under solar radiation

indicdent at the earth's surface. Of the remaining

classes, the ethers appear the most reactive and

the acetates the least reactive.

Photodissociation of Ketones Under Ambient

Conditions . Ketones play a similar role to aldehydes in that they photolyse to produce radicals which promote the oxidation of NO to N0 2 with the concomitant formation of photochemical smog. Ketones were suggested to be precursors of peroxyalkyl radicals by Purcell and Cohen (1967) who examined the photoxidation of 2-methyl-lbutane in the presence of acetone. They found

the rate of oxidation of the olefin increased as the ratio of ketone to olefin increased. The reactivity of ketones and other oxygenates under simulated atmospheric conditions has been studied

by several workers (Burnelle et al. 1966; Wilson

and Doyle, 1970; Levy and Miller, 1970; Laity

et al. 1973).

Ketone photolysis has been described by Calvert

and Pitts (1966) and the photolysis and photo- oxidation of ketones have been summarized recently

by Lande et al . (1976). The absorption spectra

of some common ketones are shown in figure 8 (Calvert and Pitts 1966). Due to uncertainty in the behavior of ketones under atmospheric

conditions, radical production is often assumed

to be 100 percent efficient. For example for MEK,

the reactions are

39
Table 10. Rate constant data for the reaction of OH radicals with other oxygen-containing organics.

Reactant

10 12 x_k

cm

3 molec~ 1 s~1 at T K Technique Reference

Ketones

Methyl ethyl ketone 3.4 ± 1.0 305 ± 2 Relative

rate Winer et al. (1976) (relative to OH + isobutene = 4.80 x 10"' 1

) 3 Methyl isobutyl ketone 14 ± 4 305 ± 2 Relative

rate Winer et al . (1976) (relative to OH + isobutene = 4.80 x 10 ") a Di isobutyl ketone 24 ± 7 305 ± 2 Relative

rate Winer et al. (1976) (relative to OH + Isobutene = 4.80 x 10~ M

) a Ketene >1. 7 ^ 295 Relative

rate

Faubel , Wagner, and Hack (1977) (relative to OH + C 3 0 2 = 1.4 x 10~ 12 Faubel et al . (1977) Ethers

CHsOCH, wl 1 3 ul« n 3 3.5 ± 0. 35 V 299 FP-RF Perry, Atkinson, and Pitts (1977) Covered T range 299-424 K obtained

A = 1.29 x 10 11 cm

3 molec~'s~' and

E = 770 ± 300 cal mol" 1 Diethyl ether 8.9 ± 1.8 305 ± 2 Relative

rate Lloyd et al . (1976) (relative to OH + Isobutene = 4.80 x 10~") a Di-n-propyl ether 16.3 ± 3. 3 305 ± 2 Relative

rate Lloyd et al . (1976) (relative to OH + Isobutene = 4.80 x 10~") a Tetrahydrofuran 13.9 ± 2.8 305 ± 2 Rg 1 3 1 i VG rate Winer et al. (1977) (relative to

OH + isobutene = 4.80 x 10~ ll

) a CH 2 = CH0CH 3 33.5 ± 3.4 299 FP-RF Perry, Atkinson, and Pitts (1977) Covered T range 299-427 K. Obtained

A = 6.10 x 10 12 cm

3 molec"'s" 1 and

E = -1015 ± 300 cal mol"' Alcohols

CH 3 0H

b 1.06 ± 0.11 292 Relative

rate Campbell, McLaughlin, and Handy (1976) (relative to OH + n-butane = 2.60 x 10" 12 )c 1.06 ± 0.11 296 FP-RA Overend and Paraskevpopulos (1978)

C 2 H 5 0H 3.3 1 0.4 909 CMC Relative

rate Campbell, McLaughlin, and Handy (1976) (relative to OH + n-butane = 2.60 x 10" 12 )C 3.74 ± 0.37 296 FP-RA Overend and Paraskevpopulos (1978) n-Propanol 4.3 ± 0.4 292 Relative

rate Campbell, McLaughlin, and Handy (1976) (relative to OH + n-butane =

C

.

oO X 10 ) L 5.33 ± 0.53 296 FP-RA Overend and Paraskevpopulos (1978) Isopropanol 6.7 ± 1.3 305 ± 2 Relative

rate Lloyd et al. (1976) (relative to OH + isobutene = 4.80 x 10" 11

) 3 5.48 ± 0.55 296 FP-RA Overend and Paraskevopoulos (1978) n-Butanol 7.6 ± 1.1 292 Relative

rate Campbell, McLaughlin, and Handy (1976) (relative to OH + n-butane = 2.60 x 10" ,2 )c CH 2 = CH - CH 2 0H 25.9 ± 3.3 440 Pulse radiolysis

Gordon and Mulac (1975)

Acetates

Methyl acetate 0.18 ± 0.05 292 Relative

rate

Campbell and Parkinson (1977) (relative

to OH + n-butane = 2.60 x 10" l2

) c Ethyl acetate 1.94 ± 0.22 292 Relative

rate

Campbell and Parkinson (1977) (re- lative to OH + n-butane = 2.60 x 10~ 12

) c

n-Propyl acetate 4.1 ± 0.8 305 ± 2 Relative

rate Winer et al . (1977) (relative to OH + isobutene = 4.80 x 10" M

) a sec-Butyl acetate 5.3 ± 1.1 305 ± 2 Relative

rate Winer et al . (1977) (relative to OH + isobutene = 4.80 x 10

_11

) a Methyl propionate 0.29 ± 0.10 292 Relati ve rate

Campbell and Parkinson (1977) (relative

to OH + n-butane = 2.60 x 10~''! ) c Ethyl propionate 1.77 ± 0.25 292 Relative

rate Campbell and Parkinson (1977) (relative

to OH + n-butane = 2.60 x 10~ 12 )c

Calculated from the Arrhenius expression of Atkinson and Pitts (1975). Osif, Simonaitis, and Heicklen (1975 ) determined rate constants relative to that for OH + CO of

k(0H + CH 3 0H)/k(0H + CO) = 0.63 ± 0.10 at 298 K and 0.98 ± 0.20 at 345 K. However, total pressures

(CH3OH + N 2 0 + CO) of 28-203 Torr were used. Since no data are available for the pressure dependence of

the OH + CO rate constant with CH 3 0H or N 2 0 as the diluent gas, no guantitative estimate of k(0H + CH 3 0H) can be made, apart from setting k(0H + CO) > 1.5 x lCfcrn'molec"^ 1 at 298 K (Perry et al . (1978)) and hence k(0H + CH3OH) >_ (0.95 ± 0.15) x 10"' 2 cm

3 molec~'s~' at 298 K. Calculated from the Arrhenius expression of Perry et al. (1976) for T = 292 K, which is also in excellent agreement with the value obtained by Campbell, Handy, and Kirby (1975).

40
Table IT. Rates of reaction of selected ox- genated hydrocarbons with the OH radical (assumed [OH] = 10 6 radicals

cm" 3

) at around 300 K.

Compound Reaction Rate (% h"

1

)

Ketones

MEK 1.2 Methyl isobutyl ketone 5.0 Diisobutyl ketone 8.6

Ethers

Dimethyl ether 1.3 Diethyl ether 3.2 Di-n-propyl ether 5.9 Tetrahydrofuran 5.0

CH 2 = CH0CH 3 12.1

Alcohols

CH3OH 0.38

C 2 H 5 0H 1.3

n-C 3 H 7 0H 1.7 I-C3H7OH 2.2 n-CH 9 0H 2.7 CH2CHCH2OH 9.3

Acetates

Methyl acetate 0.06

Ethyl acetate 0.70

n-Propyl acetate 1.5 s-Butyl acetate 1.9 Methyl propionate 0.10

Ethyl propionate 0.64

Carter et al . (1978) suggest that triplet

formation is the dominant initial process under ambient conditions and this species reacts with atmospheric 0 2 to give ethyl radicals and acetyl peroxy radicals, i.e., reaction (27). They state that the alternative route, reaction (28) is less favored thermodynamical ly. However, major uncertainty remains concerning

the photolysis of ketones under ambient conditions

including the quantun efficiency of radical production in the presence of 0 2 .

If one compared the realtive rates of aldehydes

and ketone photolysis under simulated atmospheric

conditions given by Carter et al . (1978), and uses the atmospheric aldehyde photolysis rates given earlier, one may estimate that the photolysis

rate of MEK in the atmosphere is ^ 10 percent h

1

.

7. Importance of Aldehydes and Other Oxygenates

in Modeling Atmospheric Chemistry

The smog chamber studies carried out under simulated atmospheric conditions have been mentioned earlier and adequately demonstrate the importance of aldehydes and other oxygenates in promoting photochemical smog formation. Computer

modeling studies have further emphasized the importance of these compounds. Computer results

have been shown to be significantly impacted by uncertainties in initial concentrations and photochemical parameters such as quantum yields (Niki et al., 1972; Demerjian et al., 1974; Dodge and Hecht 1975; MacCracken and Sauter, 1975; Graedel

et al. 1976; Whitten et al . 1976; Dodge and Whitten 1976; Baldwin et al . 1977; Carter et al. 1978). Thus Dodge and Hecht (1975) state that aldehyde photolysis is among the most critical reactions for quantitative photochemical smog modeling when one combines the sensitivity with

the uncertainty in the rates and mechanism of the reactions.

WAVELENGTH, A

UV absorption spectra for acetone (1), diethyl ketone (2), MEK (3), and methyl

n-butyl ketone (4) Pitts, 1966). (from Calvert and

CH,C0C 2 H 5 + hv -> CH,C0 + C 2 H ;

CH 3 + C0C 2 H 5

(27)

(28)

The situation is more complex for modeling

atmospheric conditions. Uncertainties in the photochemistry and ambient concentrations are compounded by ill-defined emission rates for aldehydes and other oxygenates. These combined

uncertainties can have a significant impact upon

model calculations. For example, figure 9 shows

the results of two runs carried out with an atmospheric trajectory model being developed by Environmental Research and Technology (ERT), under Coordinating Research Council , (CRC) funding. The photodissociation rates used in the model are those

given by Peterson (1976). This model is being

developed using data taken during the Los Angeles Reactive Pollutant Program (LARPP) carried out in 1973 (Martinez and Parker, 1976). The model partitions the hydrocarbons into five classes including separate classes for HCH0 and other

aldehydes, RCH0. It is evident from the results

shown in figure 9, that the two parameters which

we varied in this calculation can have a significant impact on the results. Of course,

each parameter must be caried separately to isolate individual effects. The choice of initial aldehyde concentrations is important in controlling

the radical concentration available to promote

N0 2 formation, but the absolute effect is modified

41
k HCHO" 2 - 0k RCHO

[ALDEH YDES]q = 3.6 pphm [FORMALDEHYDE] q = 2.5 pphm

k HCHO " k RCHO

[ALDEHYDES]q = 1.8 pphm [FORMALDEHYDE] 0

= 2.5 pphm

6.40 am 12 noon 5:20 pm LOCAL TIME (minutes)

Fig. 9. Effect of aldehyde photolysis rates and

initial concentrations on a trajectory

model run for November 5, 1973 in the Los Angeles basin.

by the source input fluxes for aldehydes.

8. Summary

The above discussions have covered the sources, ambient concentrations, radical reactions and photochemistry of aldehydes, and to a lesser extent, ketones, alcohols, ethers and esters. In general , the photolysis and reactions with the OH radical appear to be the major sinks for oxygenated hydrocarbons in the lower atmosphere. The

state of knowledge of OH radical reactions with

these oxygenates is currently adequate for modeling

purposes given the uncertainties in other areas. Consequently, a major thrust towards further refinement in these rate constants specifically

for tropospheric modeling purposes has little merit

and attention should be focused on other areas. These areas may be summarized under the general

categories below.

Ambient Concentrations . Much greater emphasis

should be placed upon obtaining concentration-time

profiles for aldehydes and ketones in the atmosphere as a function of location, season and time of day. The increase in chemical sophistication of atmospheric models, which have been validated using well characterized smog chamber

data, places increased demands on the quality and extent of ambient air quality measurements. This

is typified in the case of aldehydes and ketones

for which chemical reactions may exist in the computer model mechanism, but for which no ambient

air quality data are available either for initial conditions or to test the predicted concentrationtime behavior of these pollutants.

A greater knowledge of aldehyde and other

oxygenate emissions is also needed to form the basis of a good emission inventory for modeling

purposes

.

It is realized, of course, that the above requirements have not been met in many locations

for the common hydrocarbon classes of alkenes, alkenes and arenes and in some cases, not even for non-methane hydrocarbons.

Photooxidation . The photodissociation of aldehydes and ketones appears to be the major

depletion mechanism for these compounds in the lower atmosphere based on the calculations present

ed earlier. Although there has been significant . advances in our knowledge of formaldehyde quantum:

yields, the ambient photolysis rates of other

aldehydes, and particularly ketones, are poorly

known. Consequently, further studies of the photolysis of aldehydes and ketones as a function

j of pressure up to atmospheric, and in the presence

of 0 2 should be carried out. The uncertainty in quantum yields need to be reduced substantially

to about ± 25 percent, since the oxygenate photo

lysis steps are important in radical production.

Kinetics and Mechanism . As indicated above,

the general status of knowledge for the most

important free radical, OH, is satisfactory for modeling purposes. The mechanism of reaction should receive further attention in the areas of: HCO oxidation under ambient conditions; OH and

H0 2 addition to formaldehyde as suggested by

Horowitz, Su and Calvert (1978); the oxidation of aromatic aldehydes under ambient conditions, and the photooxidation of ketones and other oxygenates

under ambient conditions.

Finally, the kinetics of the possible HSOi, radical reactions with aldehydes and ketones should be studied to test the suggestion of Benson1 (1978). Acknowl edgments

The trajectory modeling work is being performed' under Coordinating Research Council funding. The [ author wishes to thank Doctors Roger Atkinson,

Karen Darnall and Arthur Winer of the Statewide Aii Pollution Research Center, University of Californij Riverside for helpful comments on this manuscript

and for access to results prior to publication; Professor J. G. Calvert of the Ohio State Universils and Dr. G. K. Moortgat of the Max-Planck Institute' Mainz, for receipt of results prior to publication1 }, and the Chemical Kinetics Data Center of the Natio,

al Bureau of Standards for providing a bibliograph,

on aldehyde photooxidation.

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Moortgat, G. , Private Communication, (March, 1978]

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(1978a).

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J. Amer. Chem. Soc . 93, 3570 (1971).

National Academy of Sciences Vapor Phase Organic

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Niki, H. , Private Communication, (1978).

44
j Niki, H. , Maker, P., Savage C, and Breitenbach, L., Schuck, E. A., Eye Irritation from Irradiated Presentation at the 173rd ACS National Meeting, Auto Exhaust. Air Pollution Foundation,

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in Chemistry Series No. 113, 16 (1972).

I Niki, H. , Daby, E. E., and Weinstock, B., 12th Symp. Combustion p. 277 Comb. Institute (1969).

j Niki, H., Maker, P. D., Savage, C. M. , and Breitenbach, L. P., J. Phys. Chem , in press

(1978).

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Detroit, January 9-13, 1967.

Osif, T. L. and Heicklen, J. , J. Phys. Chem . 80, 1526 (1976).

Osif, T. L. , Simonaitis, R., and Heicklen, J.,

J. Photochem 4, 233 (1975).

I Overend, R. P. and Paraskevopoulos, G. , Private Communication (1978).

Parmenter, C. S. and Noyes, W. A., Jr., J. Amer . Chem. Soc . 85, 416 (1963).

Pate, C. T, Atkinson, R., and Pitts, J. N., Jr.,

J. Environ. Sci. Health , Environ. Sci. Eng . All, 19 (1976^

1

' Penkett, S. A., Sandalls, F. J., and Lovelock,

J. E., Atoms. Env . 9, 139 (1975).

Perry, R. A., Atkinson, R., and Pitts, J. N. , Jr.,

J. Chem. Phys . 64, 5314 (1976).

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J. Chem. Phys . 67., 611 (1977).

Peterson, J. T. , Demerjian, K. L., and Schere,

I K. L., Proceedings of the International Conf.

on Photochemical Oxidant and Its Control,

Raleigh, N.C. (1977).

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Pitts, J. N., Jr., Winer, A. M. , Darnall, K. R.

Doyle, G. J., and McAfee, J. M. , Chemical

I Consequences of Air Quality Standards and of

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Smog in the Production of Photochemical

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Purcell, T. C. and Cohen, I. R. , Atmospheric

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i|Radford, H. E., Evenson, K. M. , and Howard, C. J.,

J. Chem. Phys . 60, 3178 (1974).

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Schuck, E. A. and Doyle, G. J., "Photooxidation of Hydrocarbons in Mixtures Containing

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Stephens, E. R. , Adv. in Environ. Sci. and Tech . 1, 119 (1960).

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1976/77).

Tucazon, E. C, Graham, R. A., Winer, A. M.

Easton, R. R. , Pitts, J. N., Jr., and Hanst,

P. L., Atmos. Envir ., (1978a). Tucazon, E. C, Winer, A. M. , Graham, R. A. and Pitts, J. N., Jr., Advances in Envir. Sci. and Tech . 10 in press (1978b)

Van Ham and Nieboer, H., T. NO Nieuws . 5_, 170 (1972).

Washida, N. , Martinez, R. I., and Bayes, K. D.

Z. Naturforsch . 29A, 251 (1974).

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Chemical Society Meeting, New York, (1976).

45
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(1976).

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Pollution Control Association Meeting, San Francisco, CA, October 8, 1970, 32 pp.

Summary of Session

The discussion centered on the photolysis of formaldehyde - with and without added oxygen. Heicklen began the discussion by characterizing

the photooxidation of aldehydes as the most

important unresolved problem in air pollution

chemistry. He pointed out that there appears to

be general agreement on the mechanism and quantum

yield for the photolysis of formaldehyde in the absence of oxygen, but that when oxygen is added, results from various studies suggest that the system is actually poorly understood. Of particular concern is that the results appear to differ

from laboratory to laboratory, and sometimes within

the same laboratory, as emphasized in additional

remarks by Calvert.

Warneck reported on results obtained in his laboratory (with Moortgot, Glens, and Seiler) on the uv photolysis of CH 2 0. These experiments were

unique in that low formaldehyde concentrations

were used, and both synthetic air and pure nitrogen buffer gasses were added at pressures up

to atmospheric. These results (see fig. 7 of Lloyd's review paper) appear to be in general

agreement with other recent results, within the uncomfortably large scatter of the data. At 355

nm, a pressure effect was observed, and there no longer was 1 to 1 relation between H 2 and CO production. Warneck suggested this may be due to an interaction of 0 2 with excited formaldehyde,

possibly by hydrogen abstraction.

Parkes noted that, during his study of the photooxidation of methyl radicals using a Phillips dark lamp (peak emission at 360 nm) , the photolysis

of formaldehyde was a complication. The addition of

an atmosphere of isopentane quenched the photolysis.

At shorter wavelengths, Calvert noted that they observed no quenching with up to 50 Torr of added

isobutene. At longer wavelength (^ 390 nm) the addition of 360 Torr of C0 2 resulted in quenching,:

as measured by the decrease in H 2 yield. The addition of isobutene resulted in no further

change. \ There was some discussion of the state from which formaldehyde dissociated. Ravishankara

asked if a triple state could be involved, in light of the effect of 0 2 .

Calvert replied that Ed Lee had tried to distinguish^ between triplet and singlet formal hyde by making the triplet via energy transfer

from mercury. Apparently both dissociation

processes still occur. Calvert felt that the evidence does not support triplet involvement, that the excited singlet crosses to high vibrational levels of the ground state which then

di ssociate.

i

de-it

bu

Heicklen questioned this explanation since the ; addition of an atmosphere of nitrogen resulted in;

no quenching of the formaldehyde dissociation.

Ravishankara mentioned some work by George Atkinson on dye laser photolysis of H 2 C0, monitor^

ing CO production by resonance emission.

The question of HC0 oxidation, which is direct! related to the observation of HC00H in the photo- J

oxidation of CH 2 0, was discussed by Heicklen and - Niki. In his contributed remarks, Heicklen, discussed the problems associated with the reactiqS

of HC0 with 0 2 . The direct evidence suggests thatjj! the reaction products are CO and H0 2 .

The photolysis of Cl 2 -0 2 -CH 2 0 system leads to HC00H, however, which has been interpreted as evidence for the production of HC0 3 . Niki reveal

6

that they have found peroxynitric acid by adding

N0 2 to the CI 2-O2-CH2O system, further supporting

the production of H0 2

. In addition, they have investigated the possibility that HC00H is produce

in the reaction OH + CH 2 0. They photolyzed nitroL acid and formaldehyde and found very little formicj' acid, suggesting that OH + CH 2 0 is not the source of HC00H. The reaction H02 + CH 2 0 is still considered a possibility.

The discussion ended with some remarks about

some sources of aldehydes which could become more.' important in the future. Calvert noted that the use of CH3OH as a fuel in internal combustion

engines leads to large CH 2 0 emission. Demerjian

] added that diesels produce large amounts of heavei

al dehyde.

Comments

Julian Heicklen, Department of Chemistry, The Pennsylvania State University, University Park,

Pennsylvania 16802

Moortgat et al . (presented at 13th Informal Conference on Photochemistry, Clearwater Beach,

Florida, 1978) have determined the primary photo- chemical processes in CH 2 0 photooxidation in an excellent study. At 3130 & they find the free

46
, radical process (HCO + H) to occur with <j> = 0.80

I and the molecular process (H 2 + CO) to occur with

I <|) = 0.20. However there are still some naging

I problems. I ie

I The ratio of the two paths of 4.0 is consistent

I with the value of 3.2 found by Osif (Ph.D. thesis,

Pennsylvania State University, 1976), but is

| sufficiently greater than the value of 2.1 found

j in the absence of O2 (A. Horowitz and J. G. Calvert,

j Intern. J. Chem. Kinetics , in press) to suggest

! that the primary process must be altered in the

! presence of 0 2 . This is not unexpected since the

j photochemistry does not occur from the initially

! formed excited singlet state but proceeds with an induction period (P. Avouris, W. M. Gelbart, and

M. A. El-Sayed, Chem. Rev., 77, 793, 1977). Presumably the photochemistry proceeds through a

; triplet (as in other aldehydes) or another

j intermediate (HC0H) which can be attacked by 0 2

.

j If so the formation of molecular H 2 is surprising.

A further baffling point is the results

I reported by Osif; and Horowitz and Calvert; and

I Su, Horowitz, and Calvert (13th Informal Conference

on Photochemistry, Clearwater Beach, Florida,

January, 1978) that <|){C0} could exceed 6 in the presence of 0 2 . Clearly these high yields indicate

some chain process which may be due to surface

affects

1 Finally Morrison in our laboratory has studied

j CH 2 0 photooxidation at 3130 A in a Teflon-lined

11-liter cell and finds that 0 2 quenches all CO

and H 2 formation with a half-quenching pressure of

! 4 Torr. However HC00H is produced. This result

appears to be completely at variance with all other studies.

"Julian Heicklen, Department of Chemistry, The

Pennsylvania State University, University Park,

Pennsylvania 16802

en

3 There are two problems in HCO oxidation to be

c: considered here: the rate coefficient and products

of the reaction. There are three measurements of

the rate coefficient. Washida, Martinez, and Bayes

( Z. Naturforsch 29A, 1974 and Shibuya, Ebata, Obi,

iand Tanaka ( J. Phys. Chem . 81_, 2292, 1977) reported

15.6 x 10" 12 cmVs at room temperature while Clark,

, 'Moore, and Reilly ( Int. J . Chem. Kinetics 10, 427,

1978) obtain 4.0 ± 0.8 x TO" 12 cm 3 /s. Thus there

is about a 30 percent uncertainty in this number.

,1 Of more significance is the fact that Shibuya

et al . , obtain no pressure dependence for this

rate coefficient, and conclude that there is no addition reaction of 0 2 to HCO and that H abstraction is the sole reaction path. This conforms to Hunziker's observation reported at the 12th

Informal Conference on Photochemistry (June, 1976)

(that he could not find an absorption due to HC0 3 in the HC0-0 2 reaction.

I This leaves a puzzling phenomenon since both

(Osif and Heicklen ( J. Phys. Chem . 80, 1 526, 1976)

and Niki, Maker, Savage, and Breitenbach (reported

at the 173rd American Chemistry Society Meeting,

New Orleans, March, 1977) found HC00H in the ClCH 2 0-0 2 system, which they considered as evidence

for HCO3 as a precursor. HC00H has also been seen

in other systems, as well as in polluted atmospheres

If HC00H does not come from HCO oxidation, then where does it come from?

A. R. Ravishankara, Applied Science Laboratories,

EES, Georgia Institute of Technology, Atlanta,

Georgia 30332 George Atkinson has measured CO production in

H 2 C0 flash photolysis and it seems to agree with most other data.

We have measured OH + CH 3 0H (and also OH +

C 2 H 5 0H) using flash photolysis - resonance

fluorescence. It agrees well with the indirect measurements of Campbell. So the rate constants

for the other alcohols (measured by Campbell)

are also probably correct.

Recommendations

1 . Role of Aldehydes

1.1 Kinetics . We have classified the status of

the rate data into the following caegories:

Satisfactory (S) - no need for further work

Less Satisfactory (L.S.) - further work

is desirable but not of high priority Deficient (D) - more work is needed Insufficient Data (I.D.)

1.1. (a) Hydroxyl Radical Reactions HO + RCH0 *

H 2 0 + RC0

HCH0 L.S. CH3CHO S C 2 H 5 CH0 S CH 2 = CHCH0 I.D.

C 6 H 5 CH0 S

In the case of C 2 H 5 CH0 and C 6 H 5 CH0 the rate constants are from single studies which are probably reliable, since the same authors

have also studied HCH0 and CH 3 CH0 with consistent results. It would be useful to have a second, independent determination

(say be flash photolysis experiments) to confirm the data for these aldehydes.

1.1. (b) Oxygen Atom Reactions 0(

3 P) + RCH0 -> HO + RC0

HCH0 S CH3CHO D C2H5CHO D n-C 3 H 7 CH0 I D. i-C 3 H 7 CH0 I D. CH 2 = CHCH0 L S. CH 3 CH = CHCH0 S

47
1 .1 . (c) HQ 2 Reactions

There are no experimental data available,

and the present estimated kinetic paramters

need to be substantiated by experiment.

1.1. (d) CH 3 0 Reactions

It would be useful to confirm the relative

rate data relating to the reaction CH 3 0 + CH3CHO with absolute rate data on CH 3 0 + aldehydes. Such reactions are probably

more of kinetic interest than of importance

to atmospheric modeling.

1.1. (e) N0 3 Reactions

More data are needed on this class of reactions, which might be of importance

under special conditions.

1.1. (f) HSOt Reactions

No kinetic data exist for this species, which could be important in atmospheric

chemistry.

1.1. (g) Formyl Radical Reactions

The major reaction appears to be

HCO + 0 2 * CO + HO2 for which there are reasonably consistent

room temperature data. More work is needed

on this reaction, particularly in relation

to the formation of HCOOH.

1 .2 Photolysis

1.2. (a) HCHO

The photooxidation of HCHO is tied in with

the above reaction: HCO + 0 2 > CO + H0 2 - The mechanism of HCOOH formation appears

to be dependent upon the experimental

conditions, e.g., under conditions of high

[HCHO] and low total pressure, the yields

of HCOOH are relatively high. Further

work is needed to ascertain if HCOOH is an 1 important product under atmospheric

conditions

.

The quantum yield data for the free radical fragmentation process are now in satis- factory agreement.

It might be useful to have confirmation of

the single set of cross-section data for

HCHO.

1.2. (b) CH3CHO

It would be useful to carry out further work to confirm the existing photochemical

data on CH 3 CH0.

1.2.(c) Higher Aldehydes

Photochemical studies of the higher aide- ! hydes would be useful, but are not of prime importance to atmospheric chemistry.

2. Role of Ketones

2.1 Kinetics

2.1. (a) Hydroxy! Radical Reactions

It would be useful to confirm the existing

data by further study and and at the same time to obtain information on the mechanisms of the photooxidation of the lower ketones

.

! 2.2 Photochemistry

Confirmation is needed both for the existing quantum yield measurements and for thejj] photoabsorption cross-sections, preferably

I

under atmospheric conditions.

3. Role of other Oxygenated Species

Alcohols, ethers and esters appear to be of little importance in the lower atmosphere.

43
Session III

Organic Free Radicals

ORGANIC FREE RADICALS

David M. Golden

SRI International Menlo Park, California 94025

The role of free radicals in the chemistry of the lower troposphere is reviewed. Methods

of predicting and estimating kinetic parameters are discussed with particular reference to alkoxyl radical decomposition, isomerization , and reaction with oxygen. Data needs, accuracy

and priorities are considered. Keywords: Alkoxyl; kinetics; radicals; review; troposphere.

Introduction

This paper, specifically prepared for the Workshop on Chemical Kinetic Data Needs for Modeling the Lower Troposphere, is built around

several key questions proposed to the speakers by

I the organizers

il. Why is this topic important with respect

to the chemistry of the lower troposphere?

A simplified general scheme for understanding

the chemistry of the lower troposphere is given

in figure 1. We see the role of organic free

radical chemsitry in those mechanisms, and we

are quickly led to understand that modeling of

this complex chemistry will require knowledge of

many rate constants involving organic radical

species, both aliphatic and aromatic. In fact,

we can readily see that the numbers of individual

rate constants which will be needed is enormous.

Thus, we need to be able to make reliable predictions and estimations based on a carefully

selected data base.

N02 + hv

RO + N02 R0 2 0 2

Aldehyde (Ketone) + H02

hv I J

H0 2 — * OH + N02

Fig. 1. A simplified scheme for the chemistry of the lower troposphere.

2. What is the current status of our knowledge?

This question is best addressed in terms of a framework for codification and extrapolation of kinetic data which allows us to choose the pursuit

of specific data from which we can infer the maximum amount of new information.

A slight digression to remind us of the method

involved seems in order here.

Thertnochemi stry . It is impossible to begin a discussion of the theoretical basis for critical evaluation and extrapolation of thermal rate data without first discussing methods for estimating

thermochemical quantities, such as AH£ j, AS^,

and C° T for molecules.

Group Additivity . When a sufficient data base exists, we have found [l] 1 the method of group additivity to best fit the need for accuracy and ease of operation. The basic concept and assumptions involved in the group additivity

method are as follows:

For the disproportionate reaction

RNN'R + SNN'S J RNN'S + SNN'R

any additivity approximation assumes that A$ = A$ , where $ is any molecular property, and A$

is the contribution to that property due to symmetry changes and optical isomerism. For the molecular properties of interest here, AHT * 0, AC , + 0, and AS T - S = R In K , where K

= p,T To a a a(RRNN'R)a(SNN , S)/a(RNN'S)a, SNN'R), a(X) being the symmetry number including both internal and external symmetry. An additional term for entropy

of mixing, due to the existence of optical isomers, must also be included.

l Figures in brackets indicate literature

references at the end of this paper.

51
If the molecular framework NN' is two atoms or greater, these relationships imply the additivity

of group properties, which include all nearest- neighbor interactions, since a group is defined

as an atom together with its ligands (e.g., in the group C-(H) 3 (C), the central C atom is bonded

to three H atoms and one C atom). Thus, the equation

CH3OH + CH 3 CH 2 0CH 3 J CH3CH2OH + CH3OCH3

implies the additivity of the properties of the groups C-(H) 3 (C), C-(H) 3 (0), 0-(C)(H), C-(H) 2 (C)(0) and 0-(C) 2 , if the appropriate A$ = A$

a

.

We have developed group additivity methods that permit the estimation, for many organic chemicals

in the gas phase, of heats of formation ± 1 kcal/mol, and of entropies and heat capacities to

+ 1 cal/(mol-K), from which free energies of formation can be derived to better than ± 2 kcal/

mol

.

It should be noted that entropy and heat capacity are molecular properties that can be accurately estimated under much less stringent

conditions than energy (or enthalpy). Thus, the method of bond additivity seems to work quite

well (± 1 cal/(mol-K)) for estimating the former

properties, but not at all well (± 4 kcal/mol)

for the latter.

Structural Considerations and Model Compounds . If sufficient thermochemical data is lacking for

the estimation of group properties, entropy, and heat capacity can often be adequately estimated

from structural parameters of the molecule. (Enthalpy estimates are more difficult, requiring

a better knowledge of potential functions than

are usually available). The methods of statistical thermodynamics may be used to calculate C° and

S° directly for those molecules where a

" complete vibrational assignment can be made or estimated.

Also, "reasonable" structural and vibrational frequency "corrections" to the corresponding established thermodynamic properties of "reference" compounds may be made. A suitable choice of reference compound, i.e., one similar in mass

size and structure to the unknown, assures that

the external rotational and translational entropies

and heat capacities of the reference and unknown compounds will be the same and that many of the vibrational frequencies will be similar. The basic assumption is that S° and C° difference can

be closely estimated by considering only low- frequency motions thought to be significantly

changed in the unknown. Fortunately, entropies

and heat capacities are not excessively sensitive

to the exact choice of these vibrational frequencies, and estimates of moderate accuracy

may be made with relative ease.

Kinetics . The extension of thermochemical estimation techniques to the evaluation of kinetic

data rests largely on the validity of transition

state theory.

The transition state theory expression for a thermal rate constant is:

k = (kT/h) exp[-AG°t/RT]

(The units are s _1 for first order and atm" 1 s" 1 for second order). And

AG° = AH| 00 . TASj$ c

+ <AC+> [(T-300) - Tln(T/300)].

(In the ideal gas approximation we can drop the standard state notation of AH+ and ACj). If the empirical temperature dependence is represented

by

k = AT

B exp(-C/T)

A= k/h(300)

<AC

P /R> exp[(As| 00 -<ACj>)/R]

B = (<AcJ> + R)/R

C = (AHloo - <AC|>(300))/R

k = Boltzmann's constant

h = Planck's constant

AS 0 = entropy of activation at 300 K, standard

state of 1 atm. ah| 00 = enthalpy of activation at 300 K.

<Aci> = average value of the heat capacity at

1 constant pressure of activation over the temperature range 300 T K.

If we wish to express second-order rate constants in concentration units instead of pressure units, we must multiply by RT in the appropriate units. This has the effect of writing:

k = A'T

B ' exp(-C/T)

where A' = AR and B

1 = B + 1 = (<AC+> + 2R)/R

Thus, simple "Arrhenius behavior" which will

be sufficient for lower tropospheric temperature

is characterized for first-order reaction by

ACT- = -R; (AC+ = AC? = AC^), and for second-order

reactions us i Rg concentration units by AC? = -2R

(or ACt = -R). P

In the case of simple Arrhenius behavior:

k = A exp(-B/T)

log A = log(ek<T>/h) + ASt/R;

B = (AH* + R<T>)/R

Thus, the quantities Ahft , AST-, and AC T are of interest. We apply similar methods to those already discussed with respect to thermochemistry

to view rate data in a reactional framework. These rechniques are discussed in some detail by Benson [2], but certain points are worthy of re- emphasis here.

52
We begin by classifying reactions as unimolecul.

or bimolecular. (The only termolecular processes

of interest to us will be energy transfer

controlled bimolecular processes).

Unimolecular Processes .

Simple Fission: AB > A + B

Complex Fission: Molecule > Molecule + Molecule (or radical)

Isomerization : Intramolecular atom rearrangement

Bimolecular Processes .

Direct Metathesis: A + BX + AX + B

Addition: A + Molecule •> Stable Adduct

(reverse of complex fission)

Association: A + B -> A - B (reverse of simple fission)

I The first thing to notice is that of all these

reactions, only direct metathesis reactions are

not subject to becoming energy transfer limited

at high temperatures and low pressures (i.e., in

the "fall -off" region!). This means that not

only does the so-called high pressure rate

iconstant need to be estimated or known, but the

'extent of fall -off, as well. Methods are available for making fall-off correction [3].

In hydrocarbon reactions in the troposphere,

including those of aromatic compounds, we may

expect that most direct metathesis reactions

will involve the exchange of a hydrogen atom

between larger groups. A simple, semi -empirical

prescription exists for estimating the value of

AS for these types of reactions. First, one

realizes that these values are limited between

the "loosest" possible model (A-factor equals gas

kinetic collision frequency) and the "tightest"

possible model in which R- • -H- * -R' is represented

by the molecular R-R'. Experience using data in

the 300 < T/K < 700 has taught us that generally

the AS value corresponds to a transition state

anly slightly looser than the tightest possible

value.

I Since the other two classes of bimolecular

processes are the reverse of unimolecular reactions, we may consider them in that direction.

(The equilibrium constant is either known or

'estimable). Once again using experimental results

is our guide, we note that model transition state

nhich correspond to the values of AS are generally "tight". That is, we may visualize them as ninor modifications of the reactant molecule,

usually involving some increase in rotational

entropy due to slight enlargement of certain

bonds. The dominant entropic feature is usually

(the stiffening of internal rotations as a result

Of multiple bond formation or ring formation |_2]-

| Bond scission reactions present a particular

froblem, since it is particularly difficult to

.ocate a transition state. Recent work [2,4],

oth experimental and theoretical, indicates that

|hese reactions can be modeled with a transition

I

state which becomes tighter as the temperature

rises

An example of the use of these methods to evaluate rate constants for modeling the lower troposphere is taken from Barker et al. [5]:

Alkoxyl Radical Decomposition Reactions . The decomposition reactions of alkoxyl radicals provide a good example of a family of reactions

for which an adequate number of accurate studies have been made. Most of the studies have been made on t-butoxyl radicals, but several other

radicals have been studied as well. All the studies were determinations of relative rate constants and so we returned to the original data

and recomputed it on the basis of current values

for the reference reaction rate constants.

Three different reference reactions have been

used:

0* 0N0 RiCR 2 R 3 + NO + RiCR2 R 3 (1 )

0* OH RiCR2 R 3 + (CH 3 ) 3 CH -> R!CR 2 R 3 + (CH 3 ) 3 C (2)

0' 0

I II RiCHR 3 + NO -> CiCR2 + HN0 . ( 3 ) Values chosen for ki were those obtained by Batt et al . [6], and are in good agreement with

those obtained by Golden et al . [7], the value

of k 2 chosen was that determined by Berces and Trotman-Dickenson [8]; the values chosen for k 3 were derived from disproportionation/combination

ratios and values of kj [6,9]. Other reported

data were not used because their reference reaction

rates are not sufficiently well known.

The recalculated data are presented in figures

2 through 5, and the corresponding Arrhenius

parameters are presented in table 1. The data for t-BuO are the most extensive (fig. 2), covering

nearly four orders of magnitude. The individual

sets of experimental data taken independently show

a rather wide range of Arrhenius parameters and appear to be inconsistent, but taken together, the actual data give a reasonably good straight line with parameters, log k/s" 1 = 15.1 - 16.2/0. Given the entropy change of the reaction, AH^ = 41.2 Gibbs/mole, the A-factor for the reverse

reaction is A = 1

7 ' 9 M~ 1 s"1

, a value very close

to that for trie reaction of methyl radicals with isobutene (log A = 8.0) [10]. This suggests a self-consistent method for evaluating and codifying

the limited data available for the other alkoxyl

radical reaction: choose an A-factor for the reverse reaction and find the corresponding activation energy. If this unified scheme is used, the

alkoxyl decompositions can be considered together

as a class, rather than individually.

The decomposition of an alkoxyl radical is the reverse of the addition of an alkyl radical to the carbon atom of a carbonyl group, which is analogous to alkyl radicals adding to the 2-position of a primary olefin. Since data are only available for alkyl radicals adding to the 1-position of primary olefins, the assumption was
1000/T

Fig. 2. t-BuO" + Me + acetone, y> ref - H 1 ]'' °> ref. [12]; A, ref. [13]; o, ref. [14]; dashed line, ref. [6(c)]; solid line given

by log Ms- 1

) = 15.2 - 15.9/8.

Table 1. Experimental values for RO' decomposition

rates. 3

Radical log A log E ASg log A

r log A es{ E ' Ref.

Etcr 13 7 22.1 33 4 8.2 13 7 22 1 [67] i-PrO - 16 1 20.6 37 S 8.2 14 6 17 4 [68,69]

s-BuO' 16

14

4

9 18.0

15.3 37 7 8.0 14 4 1314

9-2 [70] [65d]

t-BuO" 15 1 16.2 41 2 8.0 15 2 16 3 [71-74]

a Units: E in kcal/mol; A

r

in M-'s-i; A in s"l.

made that the A-factors for addition to both ends of an olefin double bond are the same and only

the activation energies differ. Thus, A-factors

for analogous alkyl radical plus olefin reactions

were chosen from the tables of Kerr and Parsonage

[10], corrected for any difference in reaction

path degeneracy, and applied to the alkoxyl reactions

.

Assumed A-factors for the reverse reaction,

A , are summarized along with ASn, log A . , and corresponding activation energies E in table 1. A plot of E vs AHj^ is presented in figure 6 and

gives a good straight line:

E = 12.8 + 0.71 AH° AH° > 0 (4)

= 12.8 AH° < 0 .

1000/T

Fig. 3. EtO" ^ Me + CHO. o, ref. [15], line given

by log kis- 1

) = 13.7 - 21.6/6.

This equation predicts activation energies with

an uncertainty of about ±0.5 kcal/mol. It predicts that the reverse reaction has an activation energy given by

E r

= E - AH° + RT = 13.6 - 0.29 AH° . (5) Although these equations apply to ^ 400 K where

most of the experiments were carried out, the estimated activation energies will be negligibly different at ^ 300 K.

Estimated decomposition rate constants for a number of alkoxyl radicals at 300 K and atmo- spheric pressure are presented in table 2. Fall- off corrections were obtained bv use of the

Emanual RRK Integral tables [3,19]. For the experimental data available, the estimated rates

are accurate to about a factor of two, as demo- strated by comparing estimated and observed rate constants (figs. 2-5). The differences apparent

between the estimated and experimental rate constants are due to the ±0.5 kcal/mol uncertainty in estimating the reaction activation energy and round-off errors on the A-factors.

Estimates made when no experimental data are available can be appraised by using the propagation of errors equation. Since A-factor and activation energy are usually estimated independen

tly, the uncertainty in log k can be written

a

log k = a

log A

+ CT E /6 2 (6)

54
Fig. 4. i-PrO" + Me + MeCHO. o, ref. [16]; o, ref. [17]; line given by log Ms- 1

) = 14.6 - 17.8/6.

Fig Correlation between activation energy E and enthalpy of reaction AH^.

where a-, ^, a-| A

, and are standard deviations in tog k, log A, and activation energy

respectively, and 0 = 2.303 RT. Since log A is probably uncertain by + 0.5, and the activation

energy is uncertain by about ± 1 kcal/mol, the

jtalue for

k ^ 0.88 at 300 K, and k is uncertain by about a factor of eight. This

[represents a relatively favorable case for estimations, since an adequate amount of kinetic data is

Fig. 5. s-BuO" ^ Et + MeCHO. o, ref. [18]; ref. [6(d)]; line given by log k(s- ) = 14.4 • 14.6/6.

available, and it is fairly consistent. For the reactions discussed below, very little data is available, and the realiability of the estimates

is much lower.

Alkoxyl Radical Reactions with Oxygen . The only reliable Arrhenius parameters [21] known for this class of reactions are:

CH 3 0 + 0 2 + CH3O + H0 2

log k/M

_1

s _1

= 8.5 - 4.0/9 . (7)

Rates for other members of this class can only

be estimated after making assumptions regarding

variations in A-factors and activation energies.

The concomitant uncertainties in estimating

rate constants will be relatively large since

little is known about such variations.

The A-factors for this group of reactions are expected to be similar to that for methoxy

radicals, aside from the reaction path degeneracy

(n) factor; therefore, log A can be estimated as

fol lows

log Aest

/M" s _1

= 8.0 + log n (8)

Estimates for activation energy variations are rather problematical, especially when E

fl is low. Two alternative methods can be used:

55
Table 2. Estimated RO' decomposition rates.

ASR log A log A(s"

C-COcc-co6 c-cc ccc-co6

T cc-cc

c

i c-coc

12.4 9.4

37.8 36.38.2

7.5

13.7 13.8

14.6

13.6

14.4

15.2

21 .6 19.5

17.8

19.1

14.6

15.9

k(min-'

)

0.003 2.1 y 10- 3 0.6 1.7 x 10 1

0.5 0.8

1.6 x 102 2.9 x TO 1

I0 5

10 5

HOC-CC

HOCCC-CO6 HOCCC-COH

6 (H0) 2 CCC-C0H

6 (H0) 2 CCC-C(0H) 2 6 (H0) 3 CCC-C(0H) 2 C

I HOC-CO0 0

II I CC-CC

6.8

(3.5)

f

-9.7

-9.6

-30.9

-24.5

11.4

(8.2)

f

0

38.0 8.0

36.6 7.1

39.2 7.1

39.2 6.8

37.7 6.8

37.7 6.5

37.7 7.5

40.3 7.5

14.5

13.3

13.8

13.5

13.2

12.9

13.9

14.5

17.6

(15.3) 1 19.0

12.8

20.9

(18.6)

12.8

0

2.8 x I0 3 (2.2 x 10 5

) t

2.1 x lO2

2.1 x 106

1 .0 x 106

10 5

10 5

5.2

2.6

3.6

(2.2

10° , 10 2

) f

106

^Notation: HOC-CC represents H0CH 2 CHCH 3 HOCH + HCCH 3

, etc. A-factor for analogous alkyl radical + alkene association reaction.

t

- 12.8 + 0.71 HS (kcal/mol). °FaTT-off estimated from RRK tables for 1 atm, 300 k. fRate constants for 300 K and atm air. Based on Group Additivity [2], not on experimental AH2 for propane-1,

2-diol [20].

(1) Assume E

g is constant for the entire

homologous series.

(2) Assume that an empirical relationship

that holds for other radical reactions

applies to this series as well.

Table 3. Estimates: RO- + 0 2 reactions .

Reaction log (A) E = 4.0

k (min'M

E

a

= 10.6 + 0.25

x (AH°)

k (min" 1 ) A simple empirical relationship [22] that gives

E a with an uncertainty of about ± 3 kcal mol" 1 for exothermic H, OH, and CH 3 reactions is given

by equation (9)

:

E

g =11.5,+ 0.25 (AH

R

), (9)

where AH

R

is the enthalpy of reaction. For reaction (7), equation (9) predicts E = 5 kcal mol" 1

, about 1 kcal mol 1 too high. Equation (9) has two different ways to give the proper E a for reaction (1 )

;

E

a

= 10.5 + 0.25 (AH

R

) (10)

E a

= 11.5 + 0.29 (AH

R

) . (11)

Rate constants for a number of alkoxyl radical reactions were estimated by the three methods

and are presented in table 3. Considering the large uncertainty associated with equation (2) and the low activation energies, the estimates in column III of the table are highly uncertain and may well be upper limits to the correct rate constants. Similarly, rate constants in column I of the table may be near the lower limits.

m

Ea

= 11.5 + o.

x (AH°)

k (min" 1

)

CH 3 0 + 0 2 8.5 2.0 X 10 5 2.0 X 10 s 2 0 x 10 s

EtO + 0 2 8.3 1.3 x 10 s 8.2 x 10 s

1 3 x 10 s

n-PrO + 0 2 8.3 1.3 x 10 s 8.2 x 10 s

1 l"x" 10 6

i-PrO + 0 2 8.0 6.7 x 10" 1.5 x 10 6 3 7 X 10 6

n-BuO + 02 8.3 1.3 x 10 5 3.5 x 10 s 5 8 X 10 s

s-BuO + 0 2 8.0 6.7 x 10" 1.1 x 10 6 2 2 x 10 E

Effective first-order rate constants at 300 K in air (2.1 x 10 5 ppm

The overall uncertainties for this family of reactions may be estimated as before. Log A is probably uncertain by ± 0.5 and the activation energy is probably uncertain by an average of I ± 1.5 kcal/mol. Thus, a, . ^ 1.2, and the estimated rate constant is 9 uncertain by about a factor of 16. If the activation energy is uncertain by an average of ^ ± 2.5 kcal/mol, the estimated rate constant is uncertain by a factor

of ^ 80. \

56
Alkoxyl Radical Isomerization Reactions . The importance of alkoxyl radical isomerization

reactions has been inferred from smog chamber

data [23], as well as from more qualitative

consideration [24]. The estimation of the iso- merization rates is relatively straightforward but

the estimates are somewhat uncertain, as discussed

below.

A-factors for 5-membered ring (5R) and 6-membered ring (6R) isomerizations were estimated

to be 10 11,2 s~* and lO 10-

s _1 (per H-atom),

respectively. The estimate for the 5R transition

state was made by noting that in tying up the

methyl and ethyl internal rotations, the change

in entropy is about -6.6 Gibbs/mole; subtracting

another 0.3 Gibbs/mole for the reaction coordinate,

we obtain log ATM = 11.7 for three abstractable

H-atoms. Thus, for each abstractable H-atom,

log A

5R

(per-H) = 11.2.

For the 6R transition state, a model transition

state was used. For the decomposition of ethyl

vinyl ether (EVE), log A = 11.4 at 700 K. If log

A is about the same at 300 K and S U (EVE) = 82.6 Gibbs/mole [2], then the entropy of the transition

state is 74.3 Gibbs/mole. In comparing the EVE transition state and that of n-butoxyl radical,

EVE has some doube bond character, and that of n-butoxyl will be looser by about 0.6 Gibbs/mole.

n-Butoxyl has one more hydrogen atom, worth about

0.2 Gibbs/mole, and has spin, contributing 1.4 Gibbs/mole. Adding all of these corrections

gives S

0 * = 76.5 and log A = 11.4 for three abstractable H-atoms; thus log AgR

(per H) = 10.9. The uncertainties in these estimates are probably

+ 4 Gibbs/mole and log A is uncertain by ± 1

Activation energies may be estimated from the activation energies for H-abstraction by alkoxyl radicals [25] by adding a "strain" energy of 0.5

kcal/mol for 6R reactions and 5.9 kcal/mol for

5R reactions [2]. These activation energies are rather uncertain, probably ± 2 kcal/mol. The combination of the two sources of error by the propagation of errors formula gives an estimated

uncertainty in log k of ± 1.8 at room temperature.

Thus, the rates are estimated to be uncertain by about a factor of 60.

The method for estimating these rates is summarized in table 4 and estimated rates for

several alkoxyl radicals are presented in table

5. All reactions are assumed to be at the high- pressure limit.

3. What do we need to know?

First of all, we need to test some of the preceding ideas with experiments conceived for

just that purpose. Since the ideas are based on

the transition state theory formalism, it is important to address the question of limits of

validity of transition state theory. In general,

! these testing reactions should be measured under

iconditions where isolated reactions can be [observed, as the extraction of individual rate constants from complex reacting systems is fraught

with difficulty.

Preliminary results from our own laboratory

indicate that the isomerization reaction of

Table 4. R0- isomerization reactions—estimation

procedure.

E = E (abstraction) + E (strain)

Hydrogen Abstracted E (abstraction), kcal/mol

RCH 2 -H 7.2

RCH(0H)-H 6.0

RiR2 CH-H 4.1

RiR2 R 3 C-H 4.1

RC(0H) 2 -H 4.1

a

Strain Energy

5-membered ring 5.9 kcal/mol

6-membered ring 0.5 kcal/mol

A-Factor (per abstractable H)

5-membered ring A = 10 11 ' 2 s" 1

6-membered ring A = 10 10 ' 9 s

'

Estimated.

Table 5. Estimated R0- isomerization reaction rates.

Reaction 3 logA(s _I

) E(kcal/mol) k(min"')

0CCCC * H0CCCC

0

I CCCC

OH CCCC- 11.7

H0CCCC0 * H0CCCC0H 11.2

0° H0CCCC0H * (HOhCCCCOH 11.2

0- (HOhCCCCOH (H0) 2 CCCC(0H) 2 10.9

6 (H0) 2 CCCC(0H) 2 - (H0) 3 CCCC(0H) 2 10.9

OH OH

I . I CCCC0- * CCCC0H 11.4

3.7 x 10'

8.6 X 10 3

1.9 x 10"

1.9 x 10"

2.2 x 10 5

2.2 x 10 5

3.0 x 10"

OH OH OH OH

a

I .1 I .1

Notation: CCCCO- + CCCCOH represents CH 3 CHCH 2 CH 2 0- + CH 2 CHCH 2 CH 2 0H primary alkoxyl radicals predicted by Barker

et al . [5], does indeed take place at the rate expected. In this experiment, nbuO* radicals were generated from the VLPP of nbnONO in the presence of DI. The mass spectrum of products indicated the production of both nbuOD and

DCH 2 CH 2 CH 2 OH.

There are many examples of reactions for which

rate constants have been studied, but product

studies are lacking. Thus, in reactions of OH with olefins current models must arbitrarily

decide on branching ratios. This is equally true

in aromatic systems.

In all of the above discussion of estimation

of rate data, the importance of thermochemical

values for all species has been emphasized. It

is particularly important to have a good set of values for the entropy and heat of formation

of organic free radicals.

57
Very few spectroscopic assignments exist for modest-to-large size organic free radicals and entropies (and heat capacities) have generally

been estimated by methods discussed earlier. Uncertainties arise from changes in hindered

rotation barriers and changes in skeletal bending

frequencies

.

Heats of formation of organic free radicals

have been measured by a variety of techniques.

All have been extensively presented in the literature. Common techniques and problems associated with them are:

1 . Bond scission activation energy assigned as bond strength :

R-R 1 ki R + R'

(a) Ei is a function of T.

(b) k is a function of T.

-

1 (c) Small fractional errors in slope of Arrhenius line lead to large absolute

errors in Ei. (This problem is overcome

to a large extent by use of relative rate techniques)

.

(d) ki and k are functions of P.

_

i 2. Mass spectrometric techniques requiring cycles

involving ion thermochemistry :

R R + e

S * R+

+ e

IP

AP

AP-IP = AH^

0 (R) - AhJ )0 (S)

(a) Difficulties measuring IP and AP (a whole literature)

R + e R

RH -«- R + H

EA

pK

a + AH^

AH - EA = AH?(H

+

) - Ah9(RH) + AH°(R)

a T T T (a) Difficulties measuring EA and pK . a

3. Haloqenation kinetics (principally iodination)

Measurement of the rate constant ki for

ki RH + I R + HI

and the assumption that E =1+1 kcal mol" 1 -

1

(a) Validity of the assumption

Recently several workers have suggested that

the value of AH,(R) determined by iodination

techniques are two low. This says, in effect,

that if ki is correctly measured E must be lower

than 1 ± 1 kcal/mol" 1. _1

7 9.50 " 9.30

Tj 9.10 § 8.90 £ 8.70 S 8.50

I I I I

I I I CH2 D

i k I

.fnT

/CH2 " CH2 IOJ >n1[O)

log(k 3 /M-' s -M = (9.93± 0]22) 4

2003°3

± R

1 °00

.T = 965 K

- ^

i—h—*

e

I i

i

i i i i 8.50 9.00 9.50 10.00 10.50 11.00 11.50 12.00 12.50

10" K/.T

Fig. 7. Arrhenius plot. bibenzyl as precursor;

benzyl vinylether as precursor for

benzyl radicals.

9.20

9.10 f 9.00

J" 8.90 S 8.80 2 8.70 o 8.60

8.50

8.40

"I

1

Dl

, , , 4000 ± 1000 logU^/M"' s" 1

) = (9.73 ±0.21) -

2 3Q3 RT

, , 3000 ± 1000

log(k 3 /M- 1 s" 1

) = (9.58 ±0.35) -

2 3Q3 RT—

-D + jT = 1000K

T = 635 K

g 10 11 12 13 14 15 16

10" K/T

Fig. 8. Arrhenius plot. --- methathesis reaction1 involving DI; — metathesis reaction

involving HI.

We have recently tested this possibility with

several experiments where R is allyl or benzyl radical. Figs. 7 and 8 show the results. These ; Arrhenius parameters are shown to be compatible

when a suitable transition state model is chosen

j with values of AH

f

2 98 (ally!) = 39.1 ± 1.0 (corre sponding to E_j - 2.3 kcal mol" 1

) and AH.^298

(benzyl) = 46.6 ± 1.5 (corresponding to E

= 2.5 kcal mol" 1 ). Thus we see no indication of any major problem with the halogenation technique. Work on t-butyl radical is in progress.

There are other specific elementary processes

for which the rate constants (and products) requi'«J study. It is not my intention to try to develop jl a list in this paper, but this workshop as a wholjj

might consider doing so. J

58
4. How Accurately (do we need to know whatever

it is that we need to know ?

This question cannot be answered in a general

I way. It is optimistic to think that we can know

thermochemical values for the free radicals to Ibetter than ± 1 kcal mol

1 in AHL and ± 1 cal

jmol -1 deg" 1 for S°. Errors of this size will limit possible accuracy of estimates of rate

[constants to an order of magnitude. If the data being modeled justify higher accuracy, individual

rate constants will have to be measured. In Igeneral , rate constants can be expected to be jmeasured to accuracies on the order of a factor

of two but some are hard to obtain at all.

J As far as I can tell, the limiting problem in icurrent smog modeling is as much with the data

to be modeled as the kinetic and/or thermochemical

data.

5. With respect to this particular topic what do you see as the research priorities ?

I have really addressed these previously, but

(to summarize:

' General

Tests of predictions using thermochemical

kinetics

Thermochemistry of free radicals

| Product studies

I Reactions in aromatic systems

Specific (incomplete)

Branching ratios in OH reactions with

unsaturates

Heats of formation of prototype alkyl radicals

i This Workshop should develop a list like this

:ombining all topics to eliminate overlap.

i. Are there speculative problem areas that should be given some attention ?

In the general area of organic free radical

;inetics, the question of perturbation of

| 'elementary" reactions by the formation of weakly

round complexes is one that I find perturbing.

e ''j"hus, if:

I i+ 2 A + BC + A-BC -* products

[nd k_ > k 2

kexp't = Kik 2

; 'his allows for low and/or negative activation

Inergies, since E = AEi + R 2 and AE : < 0. If

> E 2 , A > A 2 T<5r k > k 2 , thus: A . =

"i -

i ~i expt

Ai/A_ )A 2 < A .

If this type of mechanism is common, it must

be taken into account and what will be needed is

a method for recognizing and predicting such occurrences

Acknowledgment . I have profited greatly from

the work of my colleagues, Dr. John R. Barker

and Dr. Alan C. Baldwin. Conversations with

these coworkers and Dr. Dale G. Hendry have been very valuable.

Support from SRI International with some help from the Environmental Protection Agency (Contract

No. 68-02-2427) is gratefully pointed out.

Some parts of this paper are taken from previous reports and publications.

References

[I] Benson, S. W., Cruickshank, F. R.

Golden, D. M. , Haugen, G. R. , O'Neal,

H. E., Rodgers, A. S., Shaw, R. , and Walsh, R., Chem. Rev . 69, 279 (1969).

[2] Benson, S. W. , Thermochemical Kinetics , 2nd Ed. (John Wiley and Sons, Inc., New York, 1976).

[3] Golden, D. M. , Solly, R. K. , and Benson,

5. W., J. Phys. Chem . 75, 1333 (1971).

[4] Smith, G. P. and Golden, D. M. , Int.

J. Chem. Kinetics (to be published).

[5] Barker, J. R. , Benson, S. W. , Mendenhall,

G. D. , and Golden, D. M. , EPA-600/3-77-1 0

Grant No. R802288, October 1977.

[6] (a) Batt, L. , McCulloch, R. D. , and Milne, R. T., Int. J. Chem. Kinetics

6, 945 (1974).

(b) Batt, L. , McCulloch, R. D. , and Milne, R. T. , Int. J. Chem. Kinetics Symposium No. 1, 441 (1975).

(c) Batt, L. and Milne, R. T. Int. J . Chem. Kinetics 8, 59 (1976).

(d) Batt, L. and McCulloch, R. D.

Int. J. Chem. Kinetics 8, 911 (1976).

[7] Mendenhall, G. D. , Golden, D. M., and Benson, S. W., Int. J. Chem. Kinetics

7., 725 (1975).

[8] Berces, T. and Trotman-Dickenson , A. F.

J. Chem. Soc. 83, 348 (1961).

[9] Heicklen, J. and Johnston, H. S.,

J. Amer. Chem. Soc. 84, 4030 (1962).

[10] Kerr, J. A. and Parsonage, M. J., Evaluated Kinetic Data on Gas-Phase

Addition Reactions (CRC Press, Cleveland,

Ohio, 1972).

[II] Cadman, P., Trotman-Dickenson, A. F., and White, A. J., J. Chem. Soc. (A) , 2296 (1971).

59
89A , 89(a) ,

[12] Bires, F . W. , Danby, C. J., and Hinshelwood, C. M. , Proc. Roy. Soc. (London) A239 , 154 (1957T

[13] Quee, N. J. and Thynne, J. C. J., Trans. Faraday Soc. 63, 2970 (1967).

[14] McMillan, G. R. , J. Amer. Chem. Soc. 82, 2422 (1960).

[15] Leggett, C. and Thynne, J. C. J. , J. Chem. Soc. (A) , 1188 (1970).

[16] Ferguson, J. M. and Phillips, L.,

J. Chem. Soc. 87, 4416 (1965).

[17] Cox, D. L., Livermore, R. A., and Phillips, L., J. Chem. Soc. (B) , 245 (1966).

[18] East R. L. and Phillips, L. , J. Chem.

Soc. (A) , 1939 (1967).

[19] Emanuel, G. , Aerospace Report No. TR- 0200(4240-20)-5.

[20] Stull, D. R. , Westrum, E. F., Jr., and Sinke, G. C. , The Chemical Thermodynamics

of Organic Compounds (John Wiley and Sons, Inc., New York, 1969); Thermo- chemistry of Organic and Orqanometall ic Compounds (Academic Press, Inc., New York,

1970).

[21] Barker, J. R. , Benson, S. W., and Golden,

D. M., Int. J. Chem. Kinetics 9, 31 (1977).

[22] "Semenov Rule," see Laidler, K. J., Chemical Kinetics , p. 132 (McGraw-Hill, Inc. New York, 1965).

[23] Carter, W. P. L., Darnell , K. R. , Lloyd,

A. C, Winer, A. M. , and Pitts, J. N.

,

Jr., Chem. Phys. Letters 42, 22 (1976).

[24] Whitten, G. Z. and Hugo, H. H. , SAI Report EF76-126, Draft Final Report

(1976).

[25] Benson, S. W. O'Neal, H. E., Kinetics

Data on Gas-Phase Unimolecular Reactions, NSRDS-NBS 21 (U. S. Government Printing

Office, Washington, D.C., 1970).

H0CH 2 CH 2 0 <- CH 2 0H + CH 2 0

AH° = -42 ± 2 -6.1 ± 2 -26.0 AH° = 9.9 ±

S° = 76 .4 59.0 53.3 AS° = 34.9

c =

p 21 .7

Ei = 12 .8 + 0.71 (9.9 ± 3) = 19.8 ± 2

For k similar to C 2 H 5 + C 2 H l4 and log A_ = 8.0 /.log Aj = 8.0 + (34.9 - 8.35)/4.58 = 13.8

log k loo = 13.8 - 19.8/9

C - 8 ' >v6i RRK correction for fall-off s = —^ = 7 ~— £ 0.75 1 atm, 300 K

.". ki(l atm, 300 K) = 0.2 s" 1

, within a factor of

50.

0

2 II H0CH 2 CH 2 0 + 0 2

-* H0CH 2 CH + H0 2

AH° = -42 0 -75 5

log A2 = 8.3

Three methods for estimating E 2 :

3.5

4.0

AHS -28

(a) E *a

= 4.0

(b) E *b

= 10.5 +

(c) E*c

= 11.5 + (AH°)

Three estimates for k :

(a) k 2 , = 2.4 x 10 M"'s

_1 a

(b) k 2b

= 5.8 x 10 M'V 1

(c)

For a 20 percent mixture of 0 2 at 700 torr, the effective first-order rate constants are:

Appendix

Application of Thermochemical Kinetics

to the Analysis of Some Recent Data

Niki et al . ( J. Phys. Chem . 82, 135 (1978)) present data involving the analysis of a complex mechanism which leads them to the conclusion that

in the HO radical -initiated oxidation of the ethylene-NO system the radical H0CH 2 CH 2 6 cleaves

in preference to reacting with tropospheric 0 2 concentrations. This conclusion is difficult to justify, viz:

(a) ki, = 1.8 x 10 3 a

(b)

(c)

ki

b

= 4.3 x 10V 1

ki

c

= 5.2 x 10V 1

average = 3.8 x 10 3 s"

Thus, k

1 (1 atm, 300 K) £ 0.2 s" 1 and k

? (300 K) £ 4 x 10 s _1 indicate that the decomposition path- way is expected to be totally negligible, if our estimation techniques are okay. For k ls there is an uncertainty of a factor of 50, and that for k 2 is probably a factor of 20. Thus, the ratio of

k^/kj can range as shown:

60
log ^) - 4.3 ± 3

Another interesting question is raised by the

'work of Herron and Huie ( J. Amer. Chem. Soc . 99 , 15430 (1977)). They found from a study of ozonealkene reactions that they could explain their

data best by invoking the production of vibrationally excited formic acid which decomposes by

several steps, the principal one of which is

j HC00H

+

-> CO + H 2 0

i If the pathway for creating of HC00H involves

the rearrangement of the intermediate CH2OO via

0 0

\ . . /V .. II +

CH 2 00 > H 2C-0 -> 0CH 2 0 -> HC0H +

we can calculate that the internal energy in formic acid must be ca. 150 kcal mol" 1

. [•CH 2 00-] is the least stable of the above species

(AH

f ^ 48 kcal mol" 1 calculated from BDE(H-CH 2 00-

|= 93 kcal mol" 1

. Since the ring closing this

species undergoes has about 10 kcal mol 1 activation energy, the excitation relative to HCOOH

(AH f

= - 90.5 kcal mol" 1

) is ^ 150 kcal mol" 1 ].

It is possible to use quantum RRK theory to estimate the rate constant for all pathways

providing that we can write the Arrhenius

jparameters

k % A

"•' (n-m + s-1 )'. k A

(n-m): (n + s -1):

it = E/hv; m = E act/hv, s = number of oscillators

j for HCOOH; s = 9; and v (geometric mean) = 1345 cm" 1

. These calculations favor the production of OH

and HC0 radicals, but all the rate constants are

so fast that this simple calculation may not be

n ~ 150hv

*v 40 AH/kcal mol" 1 E /kcal mol" 1 log A/s" log[k/s"

HCOOH - H 2 + C0 2 - 3.6 > 50 13 < 11.7

H 2 0 + CO 6.3 > 30 12 < 10.3

0

II HCO + H 0

II H + COH

0

II HC + OH

106.6

92.6

108.9

107

93

109

14

14

16

10.3

11.1

12.3

able to discriminate properly between the first

and last reaction above. However, similar calculations for larger species do not have this difficulty. Furthermore, in the case of larger

ispecies deactivation by rollision must be taken

'into account.

Summary of Session

Parkes opened the discussion by re-iterating

Golden's contention that there is no use trying to understand the reactions of one radical in a class

in isolation from the other members of that class.

Batt followed with extensive comments on his work on alkoxy radical reactions (see contributed

comments). In response to a question by Benson as to whether the production of hot radicals by

photolysis could explain the various divergent

results, Batt replied that his experiments were carried out in an atmosphere of CF4 using 366 nm radiation for the photolysis, so that any excess energy in the radicals should be quenced rapidly.

Cox pointed out that they had done some photolysis

experiments using methyl nitrate at low concentration (^ 10 ppm) and get similar results to those

of Batt and those of Heicklen, both carried out

at higher concentrations. This suggests that hot radical production in the photolysis is not the reason for the discrepancy.

Ravishankara stated that when methyl nitrate is photolyzed in the banded region, the excess energy

is predicted to go to NO, whereas in the continuum, CH 3 0 comes off with excess energy and

NO is in the ground state.

In discussion on the rate constant for the reaction CH 3 0 + 0 2 > CH 2 0 + H0 2 , it was apparent

that not only was the value not completely agreed

upon, even the extent of the discrepancy was subject to some controversy. Much of this appeared to be due to different conceptions of what is reasonable agreement. Cox felt that,

compared to other reactions of this type, it is known rather badly. There is a discrepancy of at least a factor of three in the room temperature

value. Batt is reporting 10 9 ' 5

1 mol

-1

s

1 for the

A factor, while Golden spoke of 10 8 ' 5

, which

appeared to agree with the other data. Batt pointed out that, due to the pressure dependence

of CH3O + NO, Heicklen' s value at room temperature

could be high, bringing the A factors into better

agreement.

Tsang pointed out that, for a given class of reactions, the pre-exponential factor for decomposition is a constant, which is a valid base point for comparison.

The discussion of alkyl peroxy radicals was initiated by Benson (see contributed discussion)

who proposed a complete new mechanism for their

self reaction. Briefly, for alkyl peroxy radicals with an a-hydrogen, the reaction proceeds through

a radical disproportionate to produce a carbene peroxide radical and a hydroperoxide: 2RCH 2 0 2

•> RCH 2 0 2 H + RCH00. Heicklen asked how peroxides

(R00R), which he has observed in these systems,

are formed. Huie asked if secondary ozonides had been observed in any of these systems. Benson replied that the reaction of an aldehyde and the carbene peroxy radical to form a secondary ozonide

is slow, so it might not compete. Huie questioned

this, pointing out that work from both his laboratory and Niki's showed secondary ozonide

formation at low aldehyde concentrations.

61

L
The focus of the discussion was shifted to H0 2 , as the simplest peroxy radical. Golden, noted

that two unpublished papers suggest that its heat

of formation is 5 kcal/mol lower than presently

accepted.

Calvert brought up the question of hydration

with reference to HO2 reactions. Recent work on the recombination of H0 2 suggested that H0 2 was

hydrated. Calvert asked if other peroxy radicals

could be hydrated, or even if other classes of radicals, like alkoxy and hydroxyl , could also be hydrated.

Cox (see contributed comments) discussed his work on H0 2 recombination, which suggests the involvement of an H 2 C\ intermediate. His results

are consistent with hydration of H0 2 .

Tsang asked how much peroxy recombination

reactions affected photochemical smog. Apparently,

they are not important since as Calvert pointed

out, in an urban environment, NO was always being

pumped in. The importance of these reactions in the clean troposphere was noted by Warneck.

Tsang discussed the discrepancy in alkyl radical studies (see contributed remarks). In addition, he emphasized that the systematic

approach to radical chemistry outlined by Golden

is the only way to solve these problems.

Basco presented results on ethyl radical re- combination by flash photolysis, which is in agreement with the value reported by Parkes. Also, rate constants measured for the methyl peroxy

and ethyl peroxy radicals agree with the values

of Parkes. Calvert noted that this rate constants

also agreed with those of Parkes and Hochanadel

,

but the extinction coefficients appear to differ, which suggest the agreement in rate constants may

benfortuitious.

Basco and Parkes both stated that they have obtained a spectrum for the acetyl radical, which

should allow them to carry out kinetic measurements .

In closing, Tsang asked if spectroscopic

techniques might be useful in studying the properties of large organic radicals. It was agreed

that uv spectroscopy would not separate these

species, but infrared spectroscopy looked

promising.

Comments

L. Batt, Chemistry Department, University of Aberdeen, Aberdeen, Scotland AB9 2UE

Our work on the decomposition of alkyl nitrites

(RONO) has resulted in the determination of values

for k 2 and k 6 (tables 1 and 2 respectively):

RO + NO

RONO

RO + NO

RO

RO + 0 2 RONO

RO + NO RONO Products

R0_

H + H0 2 R0_

H + HNO

(1)

(2)

(3) (4)

(5)

R0_

H + HNO (6) These studies show that reaction (6) accounts for1 the production of nitroxyl (HNO) rather than reaction (5).

Table 1. Rate constants for the reaction

RO + NO + RONO (2)

.

R

log k 2 (M" 1 s Me

Et

i - Pr

s - Bu

t - Bu

10.1 ± 0.6

10.3 ± 0.4

10.5 ± 0.4 10.4 ± 0.4 10.5 ± 0.2

Table 2. Rate constants for the reaction

R0 + NO ^ R0_

H + HNO (6)

.

log k 6 R (M" 1 s" 1

)

Me

Et

i - Pr

s - Bu

9.3 ± 0.6

9.8 ± 0.4 9.8 ± 0.4 9.8 ± 0.4

Table 3. Arrhenius parameter for reaction (3)

Reaction

log A (s

_1

; ± 0.5

E(kcal mol

± 1

t-AmO ->- M 2 K

a + Et 14. 8 13. 8

s-BuO -> ACH b + Et 14. 9 15. 3

t-BuO M 2 Ka + Me 15 5 17. 0

i-Pro ->- ACH D + Et 14 6 17 2

EtO CH 2 0 + Me 15 0 19 8

MeO CH 2 0 + H 14 2 27 5

T Tab

?M 2 K = acetone. ACH = acetaldehyde.

Values of k 3 have been obtained by allowing

reactions (2) and (3) to compete (table 3) exceptf for R = Me. Here the value of k 3 has been obtained via a thermochemical kinetic argument. These values may be contrasted with Golden's estimated values. Figure 1 shows the first, proper and unequivocal evidence for the pressure dependence of k 3 (t-Bu0). (This means that table represents limiting values). Table 4 indicates

the variation of k 3 at 160 °C as a function of pressure and shows that k 3 is within a factor of

2 of its limiting value at a pressure of 1 atmosphere of carbon tetrafluoride. Suitable

62
0 200 400 600 800

CF4 PRESSURE (Torr)

Fig. 1. Pressure dependence of k

3 (130 °C).

| Tab! e 4. Arrhenius parameters for k 3 (t-BuO) as

a function of CFi, pressure.

log A 3 E 3 log k 3 (s

_1

) (kcal mol" 1

) (s

_1

) 11.45 11.4 5.7 no CF„

; 14.7 15.5 6.8 1 attn CH^

15.5 17.0 7.0 P > °°

conclusions may be drawn for smaller and larger

alkoxyl radicals!

Using essentially Group Additivity Rules, we

have calculated values of AH 3 for alkoxyl radicals

that appeared in Niki's schemes. We used figure 2 to determined values of E 3 (table 5). By a consideration of the reverse step we_calculate that A 3 for these radicals is ^ 10 " s" 1

. Golden came to

a similar conclusion.

30 -

A H° Kcal mol"1

Fig. 2. E 3 values determined from table 5.

Table 5. Thermochemistry for some alkoxyl radicals.

E 3 AH?

(kcal mol" 1

)

OH

CH 2 CH 2 0 -> CH 2 0H + CH 2 0

- 39.9 -4 -27.7 18 8.2

OH CH 3 CH 3 CH CH0 •+ CH 3 CH OH + CH 3 CH0

- 57.7 -13.3 -39.7 16 4.7

OH

CH 2 0 CH 2 0 + HO

- 31.8 -27.7 9.4 28.5 23.5

OH

CH 2 0 + 0 2

-* HC00H + H0 2

- 31.8 -90.5 5.0 ? -53.7

Our values for k 2 have allowed us to determined

values for k 2

' (table 6)

R0 + N0 2 R0N0 (2

1 ) We allowed reactions (2') and (4) to compete in order to determine a value for k 4 where R = Me. Table 6. Rate constants for the reaction

R0 + N0 2

•> R0N0 2 (2').

log k 2 , R (M

_1 s"

1

)

Me 9.8 ± 0.4

Et 9.9 ± 0.4

t-Bu 10.2 ± 0.4

By using dimethyl peroxide as a thermal source and

methyl nitrite as a photochemical source of methoxyl radicals, we were able to cover a temperature range from 200 °C to (Scottish) room temperature. All methoxyl radical sources now give essen- tially the same result that k h = 10 9-5 • 10" 5 /6

M

_1

s -1. It is too premature to qualify these re- sults with error limits. This result may also be contrasted to that of Golden' s. One other result

from this study is the ratio k 6 '/k 2

' <.0.l:

MeO + N0 2 CH 2 0 + H0N0 (6')

We conclude that for the alkoxyl radicals that occur in Niki's scheme, decomposition will compete

with difficulty if at all with their reaction with oxygen at room temperature. In the series R = Me,

63
Et, i-Pr, and t-Bu where we consider the two possible competing reactions of decomposition and reaction with oxygen at room temperatures:

(a) MeO reacts exclusively with oxygen

(b) t-BuO decomposes exclusively

(c) EtO and i-Pro react via both paths.

W. Tsang, Center for Thermodynamics and Molecular

Science, National Bureau of Standards, Washington,

D.C. 20234

Dr. Golden has demonstrated the important role that the thermochemistry unstable organic inter- mediate can play in providing a basis for the estimation of the rates of a variety of reactions

of importance in air pollution. It is necessary

to point out that for proper application under ambient conditions highly accurate enthalpies are necessary. A change of 1.4 kcals in activation

energy is equivalent to a factor of 10 in rate constant. In this respect some of our recent

results are especially disturbing (W. Tsang,

Int'l. Journal of Chemical Kinetics, 10, 821 (1978)). These investigations demonstrated the concordance of all existing results on the symmetrical decomposition and combination of the alkanes: n-butane * 2 ethyl, 2, 3-dimethyl butane «-

2 isopropyl and hexamethyl ethane «- 2 t-butyl

,

over a temperature range of 350-1200 K. Unfortunately, this very satisfactory agreement between

four completely disparate type of experiments

(comparative rate single pulse shock tube, radical buffer, very low pressure pyrolysis and modulation

spectroscopy) also leads to the conclusion that

the commonly accepted heats of formation of ethyl,

isopropyl and t-butyl radical are '10, 10 and 20

kJ higher than currently accepted numbers from metathesis reactions. The implication of these

results is that if there are serious questions for

such simple radical species, then, what degree

of confidence can one have for more complex

systems? Thus the utilization of estimation

schemes may be badly flawed by the uncertainty in the data base.

Sidney W. Benson, Chemistry Department, University

of Southern California, Los Angeles, CA 90007 Although bimolecular reactions of the two R0 2

" radicals do not seem to be of importance in the modeling of the tropospheric photochemistry, their self-reactions provide an important clue about an intermediate which is assuming more and more

importance, namely the carbene dioxide R'R"C00

more commonly referred to as the Criegee

zwitterion.

During the past two years, under Army sponsorship, we have been reexploring in our laboratories

at USC the reactions of interest in ignition and combustion. In these systems the self-reactions

of R0 2

' radicals assumes great importance. It has been known for some time that when R represents a tertiary carbon grouping such as Me 3 C or F 3 C, these self-reactions follow the path:

2R0 2 + 2R0 + 0 2

this overall reaction proceeds through the formation of a weakly bonded tetraoxide:

(AH

r ^ 9 kcal)

:

2R0 2 R0000R + 9 kcal

R2 0n is stable only at T < -70 °C but at least two groups of workers have studied the reversible equilibrium in solution below 200 K and made measurements of AH and AS. For R0 2 radicals where

R does not contain a tertiary C atom but has instead ct-H atoms such as R'CH 2 00* or R'R"CH00\

the equilibrium has not been observed but only a very rapid reaction leading to termination. The major products observed for such terminations are

the conjugate alcohol -ketone (or aldehyde) pair.

G. Russell some time ago proposed the following mechanism for it:

H 0 2R'CH 2 0 2

* ? R'CHzOuCHzR' «• R-C 0

H 0 \ / 0

R'CHO + 0 2 + H0CH 2 R' CH 2 R'

This is a 1,5 H transfer reaction which has never been observed in saturated molecules but only

in molecules containing at least one multiple (pi) bond. There are many objections to it and in previous publications I have discussed some of

them. Its longevity may be attributed chiefly to

its apparent ability to predict the major termination products.

In our recent reexamination of these and relatedi ozone systems, my colleague Dr. P. Nangia and I have come to the conclusion that this reaction

proceeds instead through an atypical radical disproportionate to produce the Criegee zwitterion and a hydroperoxide:

2R'CH 2 0 2

' •> R'CH 2 0 2 H + R'CHOO

We estimate that AH 2 ^ -26 kcal/mol with an activation energy of about 1 to 2 kcal/mol. This estimate is in excellent agreement with the ab initio calculations recently made by Goddard et

al . on the stability of the zwitterion. The reaction

R'CHOO - RCH0 + 0(

3 P)

is endothermic by only 6 to 8 kcal but must have

an appreciable barrier of perhaps 15 kcal in excess of this because it is a spin forbidden

process. Thus the zwitterion may have a reasonable

life time at ambient temperatures in excess of

1 second.

The zwitterion is expected to react relatively

rapidly with aldehyde to form secondary ozonides

and we estimate the mechanism to be a concerted

one with a low activation energy.

64
0-0

/ \

R'CHOO + RCHO - R' - C C - R / n y \

H 0 H

This ozonide will slowly decompose into an aldehyde R'CHO or RCHO and the isomeric, biradical

R'CH0 2 , which can rapidly and exothermical ly isomerize to the carboxylic acid.

o- /

R' - C / \

H 0-

R'C

0

0 - H

It is our feeling that this is the mechanism

for formic acid production which has been observed

in both smog chamber experiments as well as in the ambient atmosphere during smog periods. There is no kinetically acceptable way in which the precursor HCO radicals which have been supposed to

be the source of formic acid in the ambient

atmosphere can do anything but produce peroxyformic acid and that only slowly relative to

H0 2 + CO production. HC0 3 H on decomposition will not produce HC00H.

We also estimate that RCH00 can donate its weakly-bonded oxygen atom to many species such as

NO, N0 2 , R0, R0 2 , S0 2 , and possibly S0 3 . It can

also in principle react with 0 2 to form 0 3 and

RCHO. However, this reaction is spin forbidden

and may have an appreciable activation energy. Despite this it may play a role in 0 3 production

in a number of oxidation experiments which have

been reported some time ago.

The only pathways suggested so far for the zwitterions in smog episodes is from the relatively slow secondary reactions of 0 3 with olefins.

However, it does not seem unreasonable to suppose

that the exothermic reactions of a-H containing

peroxy radicals with N0 2 or with N0 3 may also

occur:

R'CH,0 5

NO;

NO;

R'CHO,

H0N0 + 14 kcal

H0N0 2 + 41 kcal

In the later portions of the smog reaction when NO

has decreased and N0 2 and 0 3 have begun to peak,

such reactions may become significant. In this

period of the overall reaction H0 2 reactions with

;R0 2

' can also produce zwitterion as well as the more famil iar R00H

It is our intention to publish all of these considerations and their related antecedents in a forthcoming publication now in preparation.

Richard A. Cox, U.K.A.E.A., Environmental and

Medical Sciences Division, A.E.R.E., Harwell, Oxfordshire 0X12 0RA, England

In connection with the problem of disproportion/

combination of peroxy-radical s we have some

'experimental information which may suggest the involvement of an H 2 0^ intermediate in the reaction

H0 2 + H02 = H 2 0 2 + 0 2 (1). The overall rate constant ki exhibits a substantial negative temperature dependence (exp(+1250/T) ) and a small pressure effect (30 percent decrease between 760 and 40 Torr). Other unpublished results from

Burrows and Thrush (Cambridge, U.K.) suggests an even lower rate constant at ^ 5 Torr. These

effects may possible be understood in terms of the formation, by the combination of 2 H0 2 radicals,

of an H2O4 molecule which can be vibrational ly relaxed by energy transfer:

H0 2 + H0 2 + H 2 0

(+M) H 2 0 2 + 0 2

If formation of products H 2 0 2 + 0 2 (as opposed

to redissociation to 2H0 2 ) from the vibrational ly relaxed H 2 0.» molecule is relatively more favored

then from H 2 0i,*, factors favoring population of

the lower vibrational levels of H 2 0i, will tend to increase the overall rate constant.

William P. L. Carter, Statewide Air Pollution

Research Center, University of California,

Riverside, CA 92521 There is inadequate data concerning the effect of non-hydrocarbon substituents on the decompositions of alkoxy radicals. Substituted alkoxy

radicals of many types are formed in polluted

tropospheric systems, and in some cases, it is uncertain whether decomposition of reaction with

0 2 predominates. As discussed in the previous

section uncertainties in the decomposition rates

of g-hydroxy-al koxy radicals present serious problems in developing models for the OH-olefinN0 system. In addition to g-hydroxy-al koxy

radicals, our detailed propene + n-butane-N0x -air

model 1 predicts formation of species of the types

0"

I R-CHCH 2 0N0 2

0

I RCH0C-R 1 , and R-CH-CH0

and it is probable that in a more complete detailed

smog model, other types of substituted alkoxy

radicals would be predicted to be formed. It is commonly assumed that decompositions of these

species predominates, but this requires experimental verification.

Carter, W. P. L., Lloyd, A. C, Sprung, J. L., and Pitts, J. N., Jr., Computer modeling of smog chamber data: progress in validation of a detailed mechanism for the photooxidation of propene and n-butane in photochemical smog, Int . J. Chem. Kinetics 11 , 45 (1979).

65
Recommendations

An understanding of the chemical kinetic behavi- our of large organic free radicals is necessary to describe the oxidation of hydrocarbons released

into the troposphere. Before we describe specific

areas where we consider research is necessary

there are some general points that should be made - points that arise from the particular complexity

of the problem.

1. The majority of the existing chemical kinetic data on the reactions of organic free radicals have been obtained from measurements of the ratios of the rates of competing processes - very often the ratio of a propagation rate constant

to a termination. It is also true that in practice

such ratios or combinations of rate constants are often the controlling parameters. It would be generally useful therefore if experimentalists

published clearly the actual ratios that they have measured as well as the rate constants that they

have deduced from them. Further, the users of rate

data, particularly modelers, should state what is needed for given situation, absolute or relative

rate constants.

2. There are too many hydrocarbons and too many resultant radicals to be able to hope to study all their potential reactions individually. This means

that we fully support the thesis of Golden in his talk that we should consider the behaviour of classes of radicals. We should make sure that any new results are consistent with observations from other members of the same class and examine deviations carefully. At the present time, this means

that the thermodynamic properties of the key radical classes such as alkyl, alkoxy and alkyl peroxy need to be firmyl reestablished after the recent upheavals. The kinetic parameters from the different types of processes have also to be measured. At the present time we are beginning

to obtain absolute rate information for the first members of the different classes. We will have to extrapolate to larger radicals that we cannot easily produce in the laboratory to estimate

firstly, rate constants for processes that are identical to those undergone by smaller radicals,

e.g., abstraction from a substrate hydrocarbon,

and secondly, to predict rate constants for processes such as alkoxy rearrangements or cleavages that the smaller species cannot them- selves undergo. The results of recent product

studies show such processes are important and the only way they will be consistently and rationally modeled will be by the establishment of a sound

and consistent data base for the radicals that

can be studied.

3. The techniques for making careful systematic measurements for individual organic free radicals

in the gas phase at temperatures of atmospheric

interest are now becoming available. The infrared

and ultraviolet absorption spectra of several larger alkyl and alkyl peroxy radicals have been detected and used for kinetic studies. The agree- ment between flash photolytic and modulation

methods has been good. So too, to a large degree,

has been the similarity for, for example, peroxy

radical combination in the gas and liquid phases. Kinetic studies using direct detection of alkoxy radicals would make a key contribution that would

help complete the stody of aliphatic radical chemistry. Oxidation chemistry can now do more

than just elucidate the admittedly complex mechan-' isms and can produce directly measured quantitati v! f

data.

4. These developments in detecting transients

have been accompanied by advances in the methods

of detecting stable products. Long path ir has been extended by using Fourier transform technique1 and the chromatography of peroxides is becoming

j more reliable. Such studies tell the paths that the radical reactions must follow. New techniques

such as laser pyrolysis when they are applied to I larger molecules will confirm whether these suggested paths are reasonable.

5. Classical determinations of radical thermo chemistry could perhaps be supplemented usefully I by structural information from the spectroscopic

information that is now appearing for larger

organic radicals.

6. The existing theory of chemical kinetics,

particularly transition-state theory and its deriatives provides a useful tool for rational izin the kinetic information (see 2). It is important

to emphasis however the degree of accuracy in energy measurements that is needed for prediction

at ambient temperatures using rate expression in the Arrhenius form.

The following specific points were raised durin1 the discussion:

1. Alkyl Radicals . The thermochemistry of ! these key initial building blocks is the subject j of controversy at the present time and must be clarified before we can proceed with confidence j with more complex radicals. There is no longer an! dispute about the order of magnitude of the rate ', constants of combination, but there is uncertainty! about the heat of formation and the structural details of the radicals.

2. Alkoxy Ra dicals. Here there is a glaring

need for a direct detection technique useful for kinetic studies. The cleavage reactions appear tq be relatively satisfactorily described but the rates of reaction with 0 2 need measurement. This means in turn consistency between measurements

against different reference reactions e.g. alkoxy

+ hydrocarbon rate constants. The rate parameters

for the recently proposed isomerizations of the large alkoxy radicals need testing preferably

directly, or if not, against an improced data base! of alkoxy measurements. Recently it has been claimed that alkoxy radicals have been detected by emission - if confirmed this could produce a break" through

.

3. Alkyl Peroxy . The key reactor here with NC has so far proved too fast for direct measurement I but in the light of the changed values for H0 2 + NO needs to be checked. The rates of radical combination that are important in lightly polluted situations are in the process of being firmly established but the mechanistic explanation of the results is speculative. Further information about the different channels in the combination is needed together with better product data.

66
4. Acetyl and Acetyl Peroxy . The absence of

( iscussion here on what are precursors on the

iroute to PAN illustrates an important area of

jtignorance. The recent detection of acetyl radicals

(should allow a start in this area and permit

Checking of the classical works in the field.

5. Information from Solution . There is a vast body of data on low temperature radical chemistry

in solution. This could form a useful basis for ideas, comparisons with gas phase work and surface

processes.

67
II
Session IV

N0X Chemistry

TROPOSPHERIC CHEMISTRY OF NITROGEN OXIDES - A SUMMARY OF THE STATUS OF CHEMICAL KINETIC DATA

Richard A. Cox

Environmental and Medical Sciences

Division, A.E.R.E,

Harwell, Oxfordshire OX 12 ORA United Kingdom

This paper is a review of the tropospheric chemistry of nitrogen oxides. The important

atmospheric reactions and the photolysis of these compounds are discussed and problem areas

are emphasized.

Key words: Nitrates; nitrites; nitrogen oxides; photolysis; reactions; review; tropospheric

chemistry.

1. Introduction

An important role for nitrogen oxides in the

i chemistry of the lower atmosphere, both from the

point of view of urban atmospheres (Leighton,

1961) and natural trace gas budgets, (Junge 1963,

Levy 1972) has been recognised for some time now. The term 'NO ' in air chemistry has usually been synonymous w^th the commonly known oxides of

nitrogen, NO and N0 2 , but recent fashion in ter- minology refers to total 'odd nitrogen' species

which includes, in addition to NO and N0 2 , the higher oxides of nitrogen, N 2 0 3 , N 2 0,>, N0 3 , N 2 0 5 and also the oxyacids of nitrogen, HONO (nitrous)

and H0N0 2 (nitric). A significant role is also

now believed to be played by peroxynitric acid,

HO2NO2.

In any model of the chemical transformations

in urban air, the chemistry of the organic

derivatives of the oxyacids, alkyl nitrites, alkyl nitrates and the peroxynitrates must be considered.

Especially important are the peroxyacyl nitrates

j (PAN's) which observational data show to be one

1 of the most important compounds in photochemical

smog.

An important atmospheric nitrogen oxide, not normally included under the terminology ' NO ' , is nitrous oxide N 2 0. Present knowledge dois not

point to a role for this oxide in the tropospheric

gaseous nitrogen cycle. However, observational

data suggests that there is a sizeable unidentified

sink for N 2 0 in the troposphere. It may be

I appropriate, therefore, to consider any chemical

1 kinetic data which might relate to this problem.

Finally, reduced nitrogen compounds, NH 3 and

its derivatives, should be mentioned, since the problem of coupling of the NH 3 and NO cycles has

been raised from time to time (Robinson and

Robbins, 1971). NH 3 undeniably plays an important

role in the aerosol and precipitation chemistry of nitrates. Relating more to the current chemical kinetic data assessment, is the problem of oxidation of NH 3 to NO (or N0 2 ).

2. Importance of NO in Atmospheric Chemistry

Nitrogen oxides and related species are important atmospheric pollutants in their own right,

e.g. the toxicity of N0 2 and the corrosive nature

of N0 2 and nitric acid toward many types of materials. However for the atmospheric chemist

it is the interaction of NO with other chemical species in the atmosphere and the resulting influence on the basic trace-gas cycles and the formation of secondary pollutants, which is of interest. It is in the solution of problems

arising from these interactions where chemical kinetics can play a role. These problems are primarily related, both in the natural and the polluted troposphere, to the photochemical oxidation of hydrocarbons.

It is now well-established that the atmospheric

oxidation of hydrocarbons and relative substances

(i.e. oxygenated and halogenated organics),

proceeds by a photochemical ly initiated free radical chain process. The chain carrying radicals

are OH, H02 and their organic analogues RO and

R0 2 . Nitric oxide, NO, is involved in this chain

process through its ability to convert the relatively inactive R0 2 radicals to active RO species via the general atom transfer reaction:

R0 2 + NO = RO + N0 2 (R = H, alkyl, etc.)

Nitrogen dioxide, on the other hand, acts as a chain terminating species through its recombination reactions with RO and R0 2 radicals, e.g.

HO + N0 2 * H0N0 2

R0 2 + N0 2 * R0 2 N0 2.

71
In the lower atmosphere H0N0 2 formation provides

an efficient sink for both N0 2 and HO species. Formation of peroxynitric acid and thi peroxyn Urates acts as a more or less temporary sink, since the products are thermally unstable and redissociate to R0 2 and N0 2 .

The other important role of N0 2 arises from its rapid photo-dissociation to give ground state atomic oxygen and nitric oxide:

N0 2 + hv(X < 400 nm) = 0(

3 P) + NO followed by

0+02 +M=03 +M .

This provides the source of ozone in photochemical

smog and also in the natural troposphere. Also

the primary source of OH radicals in the lower atmosphere, is through production of excited atomic

oxygen, 0

1 D, from the photolysis of 0 3 in the ultra-violet region at X <_ 310 nm:

0 3 + hv = 0(

X D) + 0 2 X < 310 nm

0(

1 D) + H 2 0 = 20H

Thus the involvement of N0 2 in the tropospheric

ozone budget, has a direct bearing on the average concentration of OH in the lower atmosphere and consequently on the atmospheric residence times of

a variety of trace-gases.

A model of atmospheric NO chemistry is therefore necessary to (a) formulate a realistic model of photochemical smog on which to base control strategy and (b) to provide an understanding of

the global tropospheric trace-gas cycles. The objective of such a model is to predict the total budget of NO from source to sink, and the distribution x of NO among the various chemical species in time ana in space. This distribution

depends on the nature and strength of the sources,

the chemical interactions within the atmosphere

and the role of the various sink mechanisms.

3. Cycle of N0

x Through the Lower Atmosphere

The currently accepted picture of the cycle of gaseous-nitrogen oxides through the lower atmosphere is similar to that formulated originally

by Robinson and Robbins (1971). It was based on observations and measurements of the distribution

of atmospheric NO , chemical reaction rate data, and meterological factors. Figure 1 shows a schematic illustration of the cycle.

Fig. 1. Life-cycle of nitrogen oxides in the troposphere.

The main source of NO is believed to be emission of NO from the ground, either from man- made sources, mainly combustion, or from soil processes. An additional source of NO is fixed atmospheric N 2 from lightning. There is current argument about the magnitude, both relative and absolute, of these sources.

Once in the atmosphere, chemical oxidation of

NO to N0 2 occurs rapidly, primarily through the reaction

0 3 + NO = N0 2 + 0 2 .

In daylight NO is reformed by photolysis of N0 2 , but is also oxidised by photochemical ly generated

radicals, i.e.

R0 2 + NO -> R0 + N0 2 .

Removal of N0 2 is primarily via formation of nitric acid, with alternative pathways via organic nitrates and pernitrates. Formation of peroxy- nitric acid and PANs is reversible but this can act as a sink if PANs are removed, for example,

by absorption at the ground. Additional removal of N0 2 can occur by reaction with 0 3 to give N0 3 . N0 3 is rapidly photodissociated (in daylight) but also reacts with NO (to reform N0 2 ) or with N0 2 to give N 2 0 5 . The latter reaction is reversible

and so N0 2 , N0 3 , N 2 0 5 and 0 3 can exist in equilibrium. N 2 0 5 can also be converted to HN0 3 by heterogeneous reaction with water.

The nitric acid, N 2 0 5 and organic nitrates can

all be removed from the atmosphere by absorption

at the ground (dry deposition) or by incorporation

into the precipitation elements - aerosol particles, cloud and fog droplets, which eventually

leads to rain-out.

4. Status of Chemical Kinetic Data

Accurate chemical kinetic data is clearly

required for the primary chemical processes involved in the transformation of NO to nitric

acid. Also of interest is data relating to all possible minor interactions which would influence

the basic atmospheric NO cycles, or which produce

unusual secondary pollutants in urban air. Due to the recent stimulus in the field of atmospheric

kinetics, data for some of these processes is now

rather well known. For some processes more and better data is badly needed, and these become self-evident during any detailed discussion of the data base. In the following paragraphs the status; of the data base is very briefly indicated for some specific areas of NO chemistry mentioned

above. The topics covered should not be considered

exhaustive, but rather a minimum set necessary for modelling the basic N0

x cycle.

A. Photochemical data - for N0 2 , N0 3 , N 2 0 5 , H0N0, H0N0 2 and H0 2 N0 2 \

I In order to calculate photodissociation rates for a species in the atmosphere, a knowledge of

the absorption cross-section a as a function of wavelength and the quantum yield(s) of the photo- dissociation pathway(s), <3>i, is required. This

is then combined with suitably averaged data for

the photon-flux in the atmosphere to obtain the

72
'J value' which is essentially a first order rate constant for photochemical removal of that species.

Reasonably reliable data are now available for

a and $1 for the photodissociation of N0 2 , HONO

and HONO; over the important wavelength regions

N0 2 + hv

= 0 + NO HONO + h v

= OH + NO HONO2 + hv

= OH + N0 2

I Data on absorption cross-sections and on the (various possible dissociation pathways of N0 3)

I N2O5 and HO2NO2 is much less satisfactory. Data [for NO3 is particularly important since recent

work has indicated that photodissociation can follow two different pathways, depending on wavelength :

N0 3 + hv

-y NO + 0 2 - N0 2 + 0

Only a limited amount of data is available on the absorption cross sections and dissociation quantum

I yields of the organic nitrites and nitrates.

B. Reactions of N0

x with odd-oxygen species

The reactions of NO and N0 2 with the odd oxygen

species 0 and 0 3 have long been recognized as

I important for aeronomy and reliable rate data is now available for the following reactions

0 + N0(+M) = N0 2 (+M)

0 + N0 2 = NO + 0 2 0 + N0 2 (+M) = N0 3 (+M)

0 3 + NO = N0 2 + 0 2 0 3 + N0 2 = N0 3 + 0 2

Less well known is the chemistry of N0 3 and N 2 0 5 and a number of investigations have sought to define the rate constants for the following

processes, all of which are needed for modelling

NO in the urban atmosphere:

N0 3 + NO = 2N0 2 (M) N0 3 + N0 2 = N2O5

I Due to the chemical complexity of these systems,

I there is some uncertainty in the kinetic parameters.

N0 3 may also undergo H-abstraction reactions with

organic molecules and some kinetic parameters for

these reactions have been reported. The effective

rate of the reaction of N 2 0 5 with water which is probably heterogeneous, is rather uncertain at this

time and could be of importance in the overall NO budget in the lower atmosphere. A quantitative

treatment of the rate of heterogeneous removal of

gaseous species on aerosol particles and cloud and

I fog droplets, which is acceptable to many modelers,

I has yet to be formulated. C Reaction of N0

x with odd-hydrogen species

Ifl

I The coupling of the NO and HO cycles is one

of the most important aspects of Atmospheric freeradical chemistry. The reactions of hydroxyl (HO) radicals with NO species has been widely studied

in response to problems of aeronomy, and a reasonably good data base is available here. The important reactions are:

HO + NO (+M) = HONO (+M) HO + N0 2 (+M) = H0N0 2 (+M) HO + HNO3 = H 2 0 + N0 3 HO + HONO = H 2 0 + N0 2

Note that the M dependent reactions are in the transition region between third-order and second- order kinetics at the pressures encountered in the troposphere. If the actual measurements of the rate constants as a function of pressure for

M = Air are not available, the rate constants in the transition region can be estimated from a knowledge of the third order low pressure rate constants, kjjp and the high pressure second- order rate constant k

ro

. These two rate constants

kjjj and k^ therefore comprise a minimum data

set for this type of reaction. The temperature

dependence of these association reactions is also important since in the low pressure regime they

usually exhibit a significant negative temperature

coefficient. This can be important in modelling

NO circulation in the global troposphere.

The reactions of the hydroperoxyl radical, H0 2 with NO and N0 2 are also very important. Kinetic information on H0 2 reactions is not as complete

as that for HO reactions but application of new free radical detection techniques to kinetic

studies of H0 2 has led to a significant improvement

of the reliability of the data base. Recently

new data has been reported for

H0 2 + NO + HO + N0 2

H0 2 + N0 2 (+M) t HO2NO2 + M

which allows a more quantitative appraisal of the role of these reactions. These studies have shown in particular that the reaction of H0 2 (and probably other peroxy radicals) with NO are much more rapid than had hitherto been believed, and they then assumed much greater significance in atmospheric free radical chemistry. The formation

of peroxynitric acid is reversible and associated

data for its thermal decomposition which has recently been obtained allows the role of this

'new' species to be assessed with some confidence.

D. Reactions of N0

x with organic radicals

The possible reactions of NO and N0 2 with organic radicals are numerous. However reactions with organic peroxyradicals appear to be the most significant for atmospheric chemistry, and these

are exemplified by the reactions of NO and N0 2 with peroxyacetyl radicals:

CH 3 C(0)00 + NO -> CH 3 + C0 2 + N0 2

CH 3 C(0)00 + N0 2 t CH 3 C00 2 N0 2 (PAN) These reactions govern the formation of peroxyacetyl nitrate in urban air and show clearly the competition between the chain carrying reaction
involving NO and the chain terminating step

involving N0 2 . Also since PAN can decompose back

to CH3COO2 and N0 2 , it only acts as a temporary

sink for radicals. Reasonably reliable kinetic

data for the thermal decomposition of PAN are now available and the relative rates of reaction of the peroxyacetyl radicals with NO and N0 2 are moderately well-defined.

There is little or no kinetic information concerning the analogous reactions of other organic

peroxy radicals with NO and N0 2 , or the thermal stability of the peroxynitrate species produced

in the reaction with N0 2 . Several of the peroxy- nitrates have been identified in the laboratory by infra-red spectroscopy and they may therefore

play a significant role in urban smog formation. Evaluation of the rate parameters for the analogous reactions of other organic peroxy

radicals with NO and N0 2 is necessary to define

the chemistry of the breakdown of individual organic species during atmospheric photo- oxidation.

References

Junge, C. E., Air Chemistry and Radioactivity

(Academic Press, New York, 1963).

Leighton, P. A., Photochemistry of Air

Pol 1 ution (Academic Press, New York, 1961).

Levy, H. II., Photochemistry of the lower troposphere, Planet. Space Sci . 20, 919-935,

(1972).

Robinson, E. and Robbins, R. C, Sources, abundance and fate of Gaseous Atmospheric

Pollutants-Supplement, American Petroleum

Institute Publication No. 4015, April (1971).

Summary of Session

This session was concerned with the various simple reactions of NO, N0 2 , N0 3 and N 2 0 5 , few of which are well understood. The conversion of

N0 2 into N 2 0 5 and its subsequent reaction with water to give nitric acid is of great importance. Levine noted that if the N 2 0 5 -H 2 0 rate constant

were fast enough it would play a significant role

in N0 2 removal. O'Brien discussed smog chamber

data in which the fate of NO3 depended on the substate present, suggesting aerosols may be involved in some cases in HNO3 formation. Stedman

noted that N0 2 disappearance rates at night

corresponded to the N0 2 -0 3 rate, but that the products were not able to regenerate the reactants, which suggested that HNO3 might be a product. Benson suggested an additional possible product,

pernitric acid, arising from the reaction of OH with N0 3

.

The photolysis of N0 2 is still a subject of controversy since new data (reported by Whitten)

suggest that the quantum yield is less that one around 380 to < 400 nm. It was also mentioned by

Demerjian that the same workers measured different

absorption coefficients than the NBS workers

although the latter values are still to be preferred.

Basco commented on his earlier work on the flash photolysis of N0 2 in which the 0 2 product,

arising from the secondary 0 + N0 2 reaction, was monitored. The 0 2 should be observed up to the 12th vibrational level corresponding to 46 kcal excess energy. In fact it was found up to the , 16th level which is 12 kcal higher with radiation greater than 400 nm. Whether the photoexcited

N0 2 implied by these results plays a role in atmospheric chemistry was not discussed.

The new high value of Howard for the rate constant of H0 2 + NO has been used by most modelers for the R0 2 + NO reaction. Golden pointe out that the H0 2 reaction has a negative temperature coefficient which implies a bound state. Th^ transfer of rate constants may thus be invalid. Heicklen raised the possible role of pernitrous j acid in this reaction. Hendry presented data on the pyrolysis of nitrates which support the high

| values for R0 2 + NO. Parkes also noted that in flash photolysis systems the t-butyl peroxy and

methyl peroxy radicals could never be seen in the

| presence of NO implying a very high rate constant

( ^ 10 11 cm

3 molec *s1 ). Another unresolved question was whether adducts are formed i.e.,

R0 2 + NO * R00N0 which can rearrange and cleave. Carter presented smog-chamber results which

suggest that the formation of alkyl nitrates is an important reaction. Heicklen also suggested the possibility of a reaction R0 2 + NO -* H0N0 + aldehyde involving a six member cyclic transition

state.

The reactions of R0 with NO and N0 2 were also considered. The reaction of R0 with NO leading

to the nitrate is not too important in the atmosphere (but could be in the laboratory) since

the nitrate is rapidly photolyzed. However, as Heicklen noted, part of the reaction leads to HN0 + aldehyde. It seems to be agreed that the principal fate of R0 in the atmosphere is either

isomerization, scission, or reaction with 0 2 - as discussed in the Free Radical session.

The more general problems of modeling were discussed by Dodge in terms of current problems in NO chemistry. Questions as to whether the correc photolysis rates are used and how well the models

should be expected to fit the observations were brought up but not resolved. Carter discussed

radical initiation in smog chambers, emphasizing

that it probably arises from contamination of the chamber. The real problem lies in its unpredictable nature. Solution of this problem must be given the highest priority.

Comments

Robert J. O'Brien, Patrick J. Green, and Richard

M. Doty, Department of Chemistry, Portland State University, Portland, Oregon 97207

We have been analyzing several of the UCR smog chamber experiments for their nitrogen balance.

The details of this analysis will be published in the near future and only the results will be discussed here.

For the case of a hydrocarbon which reacts only with hydroxy! radical and negligibly with other

74
free radicals or ozone we may derive an equation

for the NO balance based on the following

mechanism:

1 HC + OH -

2 N0 2 + OH

3 N0 2 + 0 3

4 N0 3 + N0 2

N 2 0 s + H 2 0

NO, + NO -

surface

loss

HNO3

N0 3 + 0 2

N2O5

2HN0 3

2N02 Under the assumption that the OH concentration

is given by

-1 . d[HC]

[0H] ' ki[HCl]

we may obtain the equation

^initial " ^0X

]

" PAN] =

k 2 r [no 2 ]

dt

kx J [HCj d[HC] + k 3 / [N02 ][0 3 ]dt (1

)

This equation assumes all N0 3 forms nitric

acid, and that the total N0 X loss (except for

PNA) is accounted for in HN0 3

.

ect

Fig. 1. Correlation of N0

X loss with reaction 2 for UCR run EC-77. T = toluene. Figure 1 shows a plot of the left hand side of eq. (1 ) vs.

:

T [N0 2 ] [T] d[T]

for UCR run EC-77 which is a toluene (T) run which

made no ozone. The slope of this line, (k 2 /ki),

500

400

300

2 200

100 -

_

1 1

1 1 1

- /

-

• -

EC -80

•

• —

-

• >r

• —

• slope = l<2'k^ =190 _

-

1 1

-

1 1 1 0 40 80 120

[no 2 i

Fig. 2.

T 0 ITI

Correlation of NO), for UCR run EC-80.

d [ T 1 ppb

loss with reaction 2 T = toluene.

is equal to 2.3 which is in good agreement with literature values for k 2 and k x . Figure 2 shows

a similar plot for EC-80, a toluene run which made

ozone. The initial slop agrees with the previous

run and the upward curvature coincides with the appearance of ozone. The deviation between the experimental data and the extrapolated straight

line is correlated with the second term on the right hand side of eq. (1) in figure 3 which is

a plot of

[NO L ]. x

J

i [N0x

] [PAN] - 1.9

[N0 2 ] TtT~ d[T]

vs. / [0 3 ][N0 2 ] dt

An excellent linear relationship is obtained.

If all N0 3 formed HN0 3 via the above mechanism we would expect the slope of figure 3 to be 2 k 3 or

.100 ppm

1 min" 1

. The actual slope is .032 ppm" 1 min" 1

.

Figure 4 gives a plot similar to figures 1 and

2 for EC-83, a run at zero relative humidity which

made 0.42 ppm 0 3 with 2 ppm initial NO . Note that there is an insignificant derivation from linearity indicating little conversion of N0 3 to HNO3 in the absence of H 2 0.

Table 1 gives a summary of 10 UCR toluene runs. The values of k 2 /ki obtained agree well with

literature values and have a standard deviation of

10 percent. The values obtained for the slope of

the curves similar to figure 3 show greater scatter

and indicate that about 1/4 to 1/3 of the N0 3 formed is converted to HN0 3 . However, the kinetic analysis is not altogether clear as to the meaning

of the linearity and value of the slope in these

plots.

75
200

Fig. 3. Correlation of upward curvature from

figure 2 with reaction 3. t = time.

1000

800 -

< 600 -

•S 400 -

200 -

Fig. 4.

T] pph

for UCR run EC-83 a dry reaction.

T = toluene.

Figure 5 shows a plot similar to figure 1 for a UCR butane run, EC-41 . Although considerable

ozone was formed, no upward curvature is observed

in this or any other butane run, indicating no conversion of N0 3 to HN0 3 . The slope of the plots

for 10 UCR butane runs are summarized in table 2 and give a value of k2 /ki of 3.7 ± 9 percent (rel. std. deviation). This value is in good agreement

with some literature values.

Table 1 Summa r> of experimental and computed data for UCR tol uene reactions

Rxn No. ECInitial NO/NO; ratio'"no'

31

PPD toluene

Maxi mum

0-, ppb

PAN ppb

Relative

humidity

%

k 2 /k, ratio

Slope of figure 3 plots ppnr'min"

77 8.9 574 276 1-2 2 40 2.3

78 2.2 1 00 230 92 40 2.6 0.026

79 4.2 1 00 980 96 15 40 2.7 0.147

80 4.2 500 1000 27 47 40 1.9 0.032

81 4.3 500 2000 313 40 2.5 0.041

82 2.0 40 1.9 0.022

83 2.1 2000 5600 420 0 1.6 0.004

84 4.9 470 970 230 70 2.1 0.019

85 4.7 50 48 2.2 0.046

86 5.1 490 1100 300 37 35 2.0 0.022

Average

Relative standard deviation

2.3

10JI

0.029

40%

- 200 -

Fig. 5. Correlation of N0X loss with reaction 2 for UCR run EC-41 . B = butane.

Table 2. Summary of experimental and computed

data for UCR butane reactions.

Rxn No. ECInitial

NO, ppb

Initial butane

ppb

Maximum

0,ppbk 2 /ki ratio

39 0 6 2.2 0 073 3.8

41 0 593 4.03 0 237 3.56

42 0 6 0.385 0 006 3.73

43 0 137 0.38 0 12 3.6

44 1 26 3.92 0 015 4.4

45 0 614 1.94 0 138 3.6

46 0 587 4.00 0 252 4.04

47 0 599 3.9 0 255 4.17

48 0 594 1.94 0 163 3.77

49 0 611 4.12 0 286 3.8

Average 3.7

Relative standard deviation

76
The very interesting tentative conclusion we

reach is that N0 3 in the presence of water is converted to HN0 3 in the toluene runs but not in the butane runs. The explanation may lie in aerosol formation which occurs with toluene but

not with butane.

For runs which form no ozone, or before ozone

buildup occurs in any run, we are able to use an analysis similar to the one in figure 1 to account

for all NO loss by the reaction OH + N0 2 + M +

HN0 3 . The

x slope of the line in this plot yields

a rate constant ratio which shows only a 10 per- cent standard deviation in the two groups of ten experiments (toluene and butane) which we have analyzed. This rate constant ratio may then be used for computer modelling of the series of

reactions with the confidence that the predicted

HC and NO loss rates will agree within ± 10 per- cent (std

x dev.) of the experimental values,

provided the correct OH concentration profile is obtained. For the case of butane the relative

HC and NO loss rates are not effected by 0 3 or

N0 3 formation and are controlled only by OH.

Donald H. Stedman, University of Michigan,

(c/o NCAR, Boulder, Colorado 80303)

It has been frequently observed in our data

and in others from rural air that N0 2 decays at night and does not come back in the morning. The

decay rate is roughly equal to the rate of N0 2 + !n0 3 . The product of the reaction is not observed

'so it is not stopping at N0 3 , presumably going on

to N 2 0 5 . When the sun comes up in the morning

the N0 2 thus lost does not reappear so the sink

cannot be photolyzed, thus it has to be going further to some product such as nitric acid.

William P. L. Carter, Statewide Air Pollution

Research Center, University of California, Riverside, California 92521

Alkyl nitrate yields observed in alkane-NO - air smog chamber runs suggest that the following

reactions

M

R0 2 + NO + RONO2

are important sources of alkyl nitrates when the

alkyl group, R, is sufficiently large [1]. No other mechanism for alkyl nitrate formation can explain the observed near-independence of the

RONO2 yield on initial NO levels, or the fact

that the possibility of H x shift isomerizations of

some alkoxy radicals do not result in significantly reduced yields of the corresponding alkyl nitrate. We consider it unlikely that the high

jyields of alkyl nitrates observed in the smog

Ichamber runs [1] could be due to heterogeneous

reactions of NO with organic products, since the measured alkyl nitrate levels do not significantly decline following NO consumption. U

If these reactions are as significant as our

mog chamber results suggest, then they would

have the effect of making larger alkanes act as radical inhibitors in photochemical smog systems,

which has significant implications concerning

their reactivity and effects on smog formation

rates.

Unfortunately, the only evidence for these reactions comes from smog chamber data, and as far as I know, our model is the only current one which includes them. Direct laboratory studies,

obtaining unambiguous mechanistic and kinetic

data concerning the reactions of R0 2 radicals with NO, are clearly required. It is particularly important that a wide variety of

R0 2 radicals be studied, in order to clearly establish substituent and size effects.

Reference

[1] DamaH, K. R., Carter, W. P. L., Winer, A. C, and Pitts, J. N . , Jr., J. Phys. Chem . 80, 1948 (1976).

Dale G. Hendry, SRI International, Menlo Park, California 94025

Peroxyacly nitrates (PANs) play an important

part in influencing the OH radical concentration

in the atmosphere. Under conditions where PANs increase in concentration they act as radical sources. This effect is associated with the facile equilibrium

ROONO2 i R00- + N0 2

where R = CH 3 C(0)0 2 , CH 3 CH 2 C(0)0 2 , etc.

Peroxyalkyl nitrates are a second type of peroxynitrate (where R = CH 3 , CH 3 CH 2 , etc.) that could also be important in affecting the OH con- centration. Richard Kenley and I are currently

studying the decomposition of these types of

compounds and, in the case of peroxy-t-butyl

nitrate, find evidence for a 20-23 kcal mol 0-N bond strength, which is considerably weaker than found for peroxyacetyl nitrate (27 kcal /mol).

When reactions 1 and -1 for peroxyalkyl nitrates are included in atmospheric models for propene, n-butane, and toluene, we find they have

no significant effect on the overall chemistry if

it is assumed that log Ai_= 16.5 s" 1

, Ei = 23 kcal/mol, log A-i = 9.0 s _1

, and E-j = 0 kcal/mol. However other reactions of peroxyalkyl nitrates

beside reactions 1 and -1 may be important. For example if the reaction

0-0

R00N02 — N

C N-0 — C = 0 + H0N02 / s // /

H 0

competed with reaction -1 then the formation of peroxyalkyl nitrates could be important radical

sinks. Additional information is needed on the chemistry of peroxyalkyl nitrates to develop

reliable atmospheric chemistry models.

77
Marcia C. Dodge, Environmental Protection Agency, Research Triangle Park, North Carolina 27711

We are encountering difficulties in our modeling efforts that appear to be related to NO chemistry. In our modeling of smog chemabe? data for the propylene-NO system, we are able to obtain respectable for all species except

0 3 . As an example, in one particular run conducted in the UCR evacuable chamber, we achieved excellent fits for NO, N0 2 , PAN, and propylene. The predicted concentrations of formaldehyde and acetaldehyde were a little low

in this run, but well within the experimental

uncertainty of the measurements. Despite the

good agreement obtained for these six species,

the peak 0 3 concentration for this particular

simulation was 42 percent too high. The experimental peak occurred at 0.38 ppm whereas

the simulated 0 3 max occurred at 0.54 ppm. Similar results are found for the other propylene/

NO runs and for several other hydrocarbon/NO

systems as wel 1

.

In our modeling we are using Howard's value of

8.1 x 10 12 cm

3 /molecule-second for the rate of reaction of H0 2 with NO. Prior to Howard's determination, we were using a rate constant of

1.4 x 10 12 cm

3 /molecule-second for this reaction.

If we repeat the propylene-NO run described above

using the old, much slower ra£e constant for the

H0 2 + NO reaction, we obtain a predicted 0 3 maximum of 0.39 ppm, in excellent agreement with

the experimental value of 0.38 ppm.

We have tried to off-set the effect of the new, fast rate constant for the H0 2 + NO reaction. Although we have varied all rate constants in the mechanism within their limits of uncertainty, we have not been able to reduce the simulated 0 3 levels without destroying the fits for the other

species. With the new H0 2 -N0 rate constant, too much NO is consumed by H0 2 late in the reaction

and 0 3 continues to build in the simulations long after the time 0 3 was observed to level off in the smog chamber. We are tempted to conclude

from this that there may be competing reactions

for H0 2 in addition to the H0 2 -N0 reaction or Howard's rate constant for the H0 2 -N0 reaction

does not apply at atmospheric pressure. Perhaps

the rate constants for the reaction of H0 2 with itself or with 0 3 , at atmospheric pressure and

in the presence of water vapor, are higher than

the values currently accepted for these reactions.

It is also conceivable that some of the H00N0 intermediate formed in the H0 2 + NO reaction may

be stabilized at atmospheric pressure so that

the effective rate of formation of OH and N0 2 is less than that measured by Howard at low pressure. It is difficult to say what the source

of discrepancy is, but the fact that many modelers are not able to handle the new, fast

H0 2 -N0 rate constant suggests that the role of

NO in smog chemistry is not yet fully understood.

William P. L. Carter, Statewide Air Pollution

Research Center, University of California, River- ! side, California 92521

Perhaps the most important single uncertainty

in NO chemistry affecting the problem of developing unambiguously validated models for tropospheric chemistry concerns initiation in smog chamber systems. As is now well known,

model simulations which assume radical initiation

only from known processes predict overall trans- formation rates in hydrocarbon-NO -air systems

far slower than those experimentafly observed in smog chambers. It is highly probable that this excess radical initiation observed in smog chamber

is due to some aspect of heterogeneous NO chemistry, since of the known species formed in these systems which can photolyze to give radicals: only nitrogen-containing species, specifically

nitrous acid or alkyl nitrites, photolize

sufficiently rapidly that contamination by

currently undetectable amounts could give the necessary rates of radical input [1]. At least

in the UCR chambers, oxygenate contamination is far lower than the levels required to give the necessary radical initiation.

This excess radical initiation in smog chamber ! systems is probably due to a chamber contamination

effect, and not to some unknown omission in the homogeneous mechanisms, or to H0N0 being

inadvertently injected along with NO . Evidence

for this was obtained in experiments x performed

at UCR employing large Teflon bags inside the black-light irradiate all glass chamber. It was observed that the overall reactivity of hydro- carbon-NO -air photolyses in a new, clean bag was ! far less £han the reactivity subsequently observed;

in the same bag after only a few smog simulation experiments were performed in it [2]. If the excess radical initiation were due to H0N0 injection or to some deficiency in the model, and I not to chamber contamination, then high reactivity

j

should have been observed in the clean, as well as

in the dirty bag.

In terms of model validation, the most serious 1 problem caused by this chamber radical source is due to the fact that it is unpredictable, and must be represented in models by some type of adjustable parameter. This means that aspects of

the model concerning radical initiation or termination cannot be unambiguously validated.

A mechanism with erroneously high radical input

in the homogeneous chemistry (such as those assuming 100 percent fragmentation to radicals in : the 0 3 -olefin system), or with erroneously low

radical termination rates, can be made to fit

the smog chamber data by suitably reducing the adjustable chamber radical input parameter, and vice-versa. These erroneous mechanisms, which appear to be "validated" by smog chamber experiments, will then give erroneous predictions in ambient air simulations, where the compensating

chamber radical parameter is removed from the model. Thus, the occurrence of this chamber

radical input phenomenon is clearly a very serious problem, and studies aimed at resolving it should

be given very high priority.

k

j

78
References

[1] Carter, W. P. L., Lloyd, A. C, Sprung,

J. L., and Pitts, J. N. , Jr., Computer

modeling of smog chamber data: Progress

in validation of a detailed mechanism

for the photooxidation of propene and n-butane in photochemical smog,

Int. J. Chem. Kinetics 1 1 , 45 (1979).

[2] Darnall, K. R., Winer, A. M. , and

Pitts, J. N., Jr., unpublished results.

W. Tsang, Center for Thermodynamics and Molecular

Science, National Bureau of Standards, Washington,

D.C. 20234

It is clear that within the forseeable future

EPA will place a great deal of dependence on smog chamber data with regard to pollution regulations

I and abatement strategies. It is therefore

I extremely disturbing to learn of the large unI certainties and ^reproducibilities in such

I experiments. One hopes that if environmental

' decisions are to have any basis in science that a vigorous program for the proper validation of

smog chamber data would be instituted as soon as possible. In particular, investigations on the nature of the surface processes, occurring on smog chamber walls should have first priority.

Recommendations

Although significant advancements have been

made recently in understanding the role of nitrogen oxides in photochemical smog, there are a number of key processes in the NO cycle for which

reliable rate data do not exist. In the session

on NO chemistry, six major areas for further

research were identified. In order of decreasing

priority, the recommended areas for investigation

are the following:

(1) Reactions of Alkylperoxy Radicals

The oxidation of NO by alkylperoxy

radical s

,

R0 2 + NO > R0 + N0 2 ,

j is an important process in photochemical smog formation. Rate constnats, however, have not been measured for this family of reactions. It is generally assumed that the rate of reaction of R0 2 radicals with NO is comparable to the rate of reaction of H0 2 radicals with NO. This assumption

may not be valid. Because of the importance of

R0 2 -N0 reactions, the temperature and pressure dependency of these reactions should be determined.

As a first priority, it is recommended that the

! rate of oxidation of NO by methylperoxy radicals

| be investigated. Methylperoxy radicals are the

i| most prevalent of the R0 2 radicals and the fate of

the CH 3 0 2 radical is important to an understanding

of the background troposphere.

In addition to the oxidation of NO by R0 2 radicals, it has been postulated that longer chain alkylperoxy radicals may add to NO to form excited

complexes that decompose to alkyl nitrates:

R0 2 + NO * (R0 2 N0)* -> R0N0 2

Since alkyl nitrate formation is radical terminating, the rearrangement shown in the above equation

could have a significant impact on smog chemistry.

It is important, therefore, to determine the" extent to which this chain-terminating reaction

can occur.

Another process of potential importance in the polluted atmosphere is the reaction of alkylperoxy

radicals with N02 to form peroxynitrates,

R0 2 + NO j R0 2 N0 2 ,

and the subsequent decomposition of the nitrates.

If this class of reactions is analogous to the

HO2 + N0 2 reaction, the alkyl peroxynitrates would

be too short-lived to be of importance in smog chemistry. They could, however, play a role at higher altitudes and at low temperatures and are worthy of study for this reason.

The possibility of the alternative channel of decomposi tion

,

R0 2 N0 2 + RCHO + H0N0 2 , should also be investigated. Decomposition to nitric acid could be important even if this route occurs to only a small extent.

To model the behavior of the alkyl peroxynitrates

(and H0 2 N0 2 as well), information is needed on the absorption cross-sections and the products arising

from the photodissociation of these nitrates. The photochemistry of these species, however, is expected to be of greater importance in the stratosphere than in the lower troposphere.

(2) Chemistry of Peroxyacyl Nitrates

It was recently determined that PAN undergoes rapid thermal decomposition:

PAN -> CH 3 C(0)0 2 + N0 2 .

The rate of this decomposition as a function of temperature is reasonably wel 1-known. Reliable

kinetic data are also available on the relative

rates of reaction of the peroxyacetyl radical with

NO and N0 2 ,

CH 3 C(0)0 2 + NO * CH 3 + C0 2 + N0 2

CH 3 C(0)0 2 + NO -> PAN

Rate data, however, do not exist for the analogous

reactions of the other peroxyacyl radicals. Atmospheric observations indicate that the PAN-type

compounds are the most stable of the organic peroxy- nitrates and, therefore, merit study. As a first order of priority, the thermal stability of the higher analogs of PAN should be determined. If these compounds are sufficiently stable, the relative rates of the reactions of the corresponding peroxyacyl radicals with NO and N0 2 should be measured. The PAN-type compounds recommended for study are peroxypropionylnitrate (PPN) and peroxybenzoyl nitrate (PBzN), important for understanding
the fate of aromatics in smog.

(3) Reactions of OH and H0 2 Radicals with NO

and N0 2

and N0 2 , The reactions of hydroxyl radicals with NO

OH + NO + (M) > HONO + (M)

OH + N0 2 + (M) + HONO2 + (M),

have been studied by a number of investigators.

Uncertainties, however, in the rates of these reactions under atmospheric conditions still exist. These pressure-dependent reactions are in the changeover region between third-order and second- order kinetics at atmospheric pressure. Because of this pressure effect, additional studies of these reactions are warranted.

New rate data have been reported recently for

the reaction

H0 2 + NO -> OH + N0 2 The new rate constant is significantly higher than previously determined values for the rate of this reaction. The faster rate constant has a substantial impact on the predictions of photochemical

models. Direct determinations of the H0 2 -N0 rate constant have been made only at reduced pressure. Because of the significant impact of this reaction

on model calculations of smog formation, the effect

of pressure on the reaction rate should be determined. In addition, the effect of water

vapor on the rate of this important reaction

should be elucidated.

(4) N0 3 Chemistry

The reactions of N0 3 have been extensively

studied by Johnston. Recently he corrected his previous rate data for the reactions,

N0 3 + NO -> 2N0 2

N0 3 + N0 2 + N2O5

An independent confirmation of the pressure and temperature dependency of these reactions is recommended.

A confirmation of the recently published absorption cross-section and quantum yields for N0 3 photolysis is also desirable. Two pathways have

been identified for this photodissociation:

N0 3 + hv N0 3 + hv

NO + 0 2 N0 2 + 0

The significance of these processes in the perturbed troposphere appears to be minimal; however, N0 3 dissociation could be of importance to an under- standing of the natural troposphere.

(5) Reactions of Alkoxy Radicals

Alkoxy radicals in the lower troposphere

undergo bimolecular reaction with 0 2 , NO, and N0 2 , To assess the importance of the alkoxy-NO

reactions, it is necessary to know the rate of reaction of RO radicals with NO and NO2 relative

to the rate of reaction with 0 2 . Such relative

rate data do not exist.

Reaction of alkoxy radicals with N0 2 can proce< by two pathways:

RO + N0 2 RO + N0 2

RONO2 RCHO + HONO

The rates of the addition and abstraction reactior for the various alkoxy radicals are reasonably

well-known. Rate constants for the alkoxy-N0 2 reactions relative to 0 2 , however, are less welldefined.

Reaction of alkoxy radicals with NO also can proceed by two pathways:

RO + NO ^ RONO

RO + NO -> RCHO + HNO

It is generally assumed that the alkyl nitrites rapidly photolyze and, therefore, do not serve as

j

effective sinks for alkoxy radicals. The rate of photolysis, however, is uncertain and merits stud>

Only one determination has been made of the relative rate of the addition reaction versus the abstraction reaction to form an aldehyde (or a ketone) and HNO. An independent confirmation of j the importance of the abstraction pathway is recommended.

(6) Heterogeneous Reactions

The assumption is generally made that heterogeneous processes are unimportant in the atmosphere. Little quantitative information is available to support this statement. As an example, the rate of removal of N2O5 by water in the lower troposphere,

N2O5 + H 2 0 + 2HN0 3

is uncertain. This heterogeneous reaction could b significant and deserves additional study. The

rate of heterogeneous removal of other gaseous species by aerosols or fog droplets could also be significant and merits attention. An estimate of

the dry deposition velocities of various reactive

NO species is also needed to assess the significance of such removal processes in the lower troposphere.

In addition to characterizing the role of heterogeneous processes in the atmosphere, it is also necessary to determine the degree to which heterogeneous processes affect the results of smog chamber studies. Kinetic mechanisms for photo80
chemical smog are tested primarily using data

, collected in smog chambers and, therefore, it is

I important to fully characterize surface effects

and other chamber-related phenomena. These phenomena include the heterogeneous formation of HN0 3 and HONO within chambers and the absorption

and desorption of reaction species from the chamber walls. One common problem encountered in modeling chamber data is that it is difficult to reproduce the observed initial rate of hydrocarbon

and NO disappearance. The very rapid initial decay of these species in smog chambers suggests

! that there is a nonhomogeneous source of free

1 radicals present at the onset of irradiation. It

is possible that radicals may be produced from

I contaminants on the chamber walls or they may arise

from the photolysis of nitrous acid. (There is evidence to suggest that HONO may form during

loading of smog chambers). It is important to the modeling effort to characterize this radical initiation process.

Summary: A number of the key processes suggested for study involve organic peroxy radicals." It is not recommended that kinetic studies be conducted on every member of each family of reactions. Only enough members of each class of reactions should be studied to establish a representative

data base. This data base should then be used

to generalize rates for the other members of the series using established thermochemical estimation

techniques.

81

Session V

Aromatics

REACTIONS OF AROMATIC COMPOUNDS IN THE ATMOSPHERE

Dale G. Hendry

Physical Organic Chemistry Group

Chemistry Laboratory

SRI International Menlo Park, California 94025

This paper is a review of the tropospheric chemistry of aromatic compounds. The reactivity of aromatic compounds is discussed and rate constants for their reactions

with OH are tabulated. The reaction mechanisms are discussed in detail.

Key words: Aromatics; free radicals; mechanism; reactions; tropospheric chemistry.

1. Introduction

Single-ring aromatic compounds make up a high

proportion of the carbon found in the urban

atmosphere. Table 1 summarizes data reported on

the composition of the atmosphere in Los Angeles,

California [l] 1

, and Manhattan, New York [2]. From these studies we see that benzene and alkylTable 1. Atmospheric concentrations of single-ring aromatic hydrocarbons.

Concentration, ppm (ppmC)

Los Angeles Manhattan

Compound (1973) (1969)

C 6 H 6 0 008 (0 048) 0.004 a (0.022)

a

CH 3 -C 6 H 5 0 020 (0 140) 0.013 (0.094)

(CH 3 )2-C 6 H4 0 020 (0 160) 0.012 (0.094)

CH 3 CH 2 -C 6 H 5 0 004 (0 032)

n-C 3 H 5 -C 6 H 5 0 001 (o 009) 0.083

s-Ci+ Hg-CkH5 0 005 (0 050) E Aromatic 0 058 (0 439) (0.294) E Alkane-CH 4 0 259 (1. 15) (0.371

)

b E Al kene 0. 087 (0. 271 ) (0.096)

2. 1 (2. 1)

0. 038 (0. 076) 0.022 (0.044) EE 0. 404 (1. 86) (0. 761

)

b

Estimated

.

b Not included ca. 0.038 ppmC ethane, ca . 0.043

ppmC propane and ca. 0 022 ppmC benzene

.

Figures in brackets indicate literature references

at the end of this paper.

substituted benzene compose roughly 24 carbon per- cent of the total hydrocarbon in Los Angeles and about 37 carbon percent in Manhattan.

The aromatic compounds in the atmosphere come from gasoline [3], in which they are used to enhance the octane rating. Gasoline itself is composed of 30 to 40 percent aromatic hydrocarbons

and apparoximately 6 to 8 percent toluene. Hydro- carbons emitted in automobile exhaust are composed

of 6 to 8 percent toluene.

We have known for some time from smog chamber reactivity studies that the al kyl -substituted

benzenes are reactive in promoting oxidation of

NO to N0 2 and formation of ozone [4]. However, only in the last few years has an effort been made

to understand specifically how these compounds

react. This effort has been very productive,

largely because it builds on an existing background

of moderately well understood smog chemistry of

the alkanes and alkenes. The total conversion of

the aromatics to H 2 0, CO, and C0 2 is a complex

process, of which we understand only the initial steps

.

2. Initial Reactions of Alkylbenzenes

Table 2 summarizes the possible reactions of toluene, a representative aromatic hydrocarbon,

with the oxidizing species known to be present in the atmosphere. Best values of rate constants and approximate concentrations are included for estimating the rate of loss of toluene by the various processes. The data in table 2 show clearly that

the only important reaction of toluene in the atmosphere is with OH. The contribution of the reactions with 0 atom and 0 3 are about 10" 14 and

10 3 of that of OH reaction. The reactions of R0 2

* proceed extremely slowly and can account for only

10 8 of the total consumption of toluene.

Rate constants for the reaction of OH with

various alkyl benzenes are summarized in table 3. We are fortunate to have several techniques for measuring the rate constants for reaction of OH with aromatic hydrocarbons. The agreement between

85
Table 2. Reactions of toluene in the atmosphere.

Concentration k,_ _ Reactant molec cm" 3 (ppm) cn^molec'V 1 tol , s

OH [5-7] 5.0 x 10 6 (2 x 10" 7

) 6.4 x 10" 12 3.1 x 10 4

0

[8] 2.5 x 10 4

(1 x 10" 9) 7.5 x 10" 14 5.3 x 10 8

°3

[9] 1.5 x 10 12 (6 x 10" 2

) 3.4 x 10 22 2.0 x 10 9

R0 2

- (H0 2 -) [10] 2.5 x 10 9 (1 x 10" 1*) 1.7 x 10 22 2.3 x 10 12

Table 3. Reported rate constants for reaction of

OH plus aromatic hydrocarbon.

Compound

Perry

[7] Hansen[6] Doyle[11] Davis[5]

Benzene 1 20 1 24 < 3 8 1.59

Tol uene 6 40 5 78 4 2 6.11

o_- Xylene 14 3 15 3 12 8 12.4

m- Xylene 24 0 23 6 23 2 20.5

p_- Xylene 15 3 12 2 12 3 10.5

1 ,2,3-Mesitylene 33 3 25 4 23 0

1 ,2,4-Mesitylene 40 0 33 5 33 0

1 ,3,5-Mesitylene 62 4 47 2 52 0

the reported values is very good, and we have confidence in these numbers. Two basic techniques

have been used: measurement of the decay of a pulse-generated OH concentration by resonance fluorescence [5-7] and determination of the rate of disappearance of the hydrocarbon relative to

a standard hydrocarbon under conditions where OH

is the sole reactive species [11].

3. Products of OH-Aromatic Reactions

Two reactive pathways are expected for the reaction of aromatics with OH radical. For toluene, these two pathways are H-atom transfer

from the methyl group (reaction 1) and addition

to the ring (reaction 2).

From the pressure dependence at 25 °C, Davis

et al . [5] suggested that ki/(ki + k 2 ) was less than 0.5. Perry et al . [7] found that the toluene-OH reaction, as well as other aromatic-OH

reactions, was strongly temperature dependent. In

fact, at higher temperatures the apparent first- order rate constants were found to be lower than

the room temperature value because of the rever- sibility of reaction (2). From extrapolation of

the data at high temperature (where only reaction

1 is important) to lower temperatures, the ratio

k i / ( k i + k 2 ) at 25 °C could be estimated. For toluene, Perry et al . postulate ki/ki + k 2 ) = 0.16 + 0-07 jne relatively large uncertainty

arises from the uncertainties associated with extrapolating the high temperature data to determine the value of ki at 25 °C.

In our laboratory we have been investigating

the products of reaction of aromatic hydrocarbons

and OH in a discharge flow system [12,13]. The products were collected in cold traps and on sol i

c

adsorbents. The product distributions were determined as a function of hydrocarbon, N0 2 , and

0 2 pressures. Table 4 summarizes some of the dat; obtained as a function of N0 2 pressures. The fraction of products resulting from reaction (1)

is a measure of ki / ( k i + k 2 ) and remains constant

over the range of conditions. For toluene we obtain 0.15 +^0.02, which agrees very well with

the best value reported by Perry et al. [7]. The

k i/ ( k i + k 2 ) values for various aromatic hydro- carbons obtained by these two methods are summarized in table 5.

Table 4. Product distribution for the reaction

of toluene plus 0Ha .

N0 2 , 10" 14 molec/cm 3

Products 7J_ 1 .04 l^39_ 1.75

C 5 H 5 CH0 11 .7 9.9 9 .7 10.4

C 6 H 5 CH 2 0H 3 .5 3.3 4 A 4.6 m-NO^H^^ 33 .3 37.6 39 .9 47.3 o-HOC^CH;, 40 .3 37.3 35 .8 29.0 p-HOCgH^CH? 6 .4 6.8 6 .7 5.5

m-H0C 5 H lt CH 3 4 .3 4.8 3 .2 2.9

H0C 6 H 4 CH 3 51 .0 48.9 45 .7 37.4

CH 3 C 6 H 3 0 2 0 .4 0.3 0 .3 0.7

C 6 H 5 CH0 +

C 6 H 5 CH 2 0H/

Total products 15 .2 13.2 14 1 14.6

a 0xygen: 9.7 x lo 16 molec/cm 3

; toluene: 3 x 10 1 molec/cm 3

; total pressure: 8.8 Torr.

We find m-nitrotol uene to be a major product;

however, the concentration varies with the 0 2 /N0 2 ratio. Thus the intermediate formed in reaction^

(2) appears to react by two parallel pathways.

86
o

*

According to this mechanism, the relative amounts of the sum of the cresols compared with m-nitrotoluene are:

[m-nitrotoluene] k L [N0 2 ] [cresols] k 3 InoT

Analyzing our data according to this expression

gives k lt /k 3 = 4 x 10 3

.

Using this value of U/k 3 , we estimated the percent yield of m-toluene as a function of N0 2 the atmosphere, as shown in table 6. Thus, at the N0 2 concentrations generally used in smog chamber experiments, m-nitrotoluene can account

for 1 to 20 percent of the toluene. At concentra

tions of N0 2 found in the atmosphere, however,

significantly less than 1 percent toluene will give m-nitrotoluene.

Table 5. Reaction of aromatic hydrocarbons

with OH.

in

k 1 /(k 1 + k ? )

Hydrocarbon This work Perry [7] Benzene < 0.05 0.01-0.13

Toluene 0.15 ± 0.02 0.01-0.23

p-Xylene 0.15 ± 0.02 0.04-0.14

Mesitylene 0.021 ± 0.006 0.01-0.04

Atmospheric

products

NO? < 1 ppm

100% Phenol

15% Benzaldehyde, 85%

cresol isomers

15% Methyl - benzaldehyde, 85% 2,5-dimethyl phenol

2% Dimethyl - benzaldehyde,

98% tri methyl

phenol

Table 5 also lists the products expected to be formed under atmospheric conditions. These

estimates are corrected for the N0 2 /0 2 ratio and

the high radical concentration. Thus, for the intermediates formed in reactions (1) and (2), the reactions important in the atmosphere are reactions (3) and (5).

Coj

^ 0 fo

In addition to the meta -nitrotol uene ortho - and para -isomers have also been reported in smog chamber experiments [14,15]. These isomers

could potentially be formed from the meta -OH-adduct

of toluene in sequences similar to reaction (4). However, since very little m-cresol is formed,

this route does not seem reasonable. In many

cases, we have observed NO to be an effective

nitrating agent upon condensing our reaction

Table 6. Calculated yields of m-nitroluene

as a function N02 concentration

in the atmosphere.

N02

10" 12 cn^molec" 1 (ppm) m-Nitrotol uene, % 1.0 (0.04) 0.07 3.0 (0.12) 0.2 10.0 (0.40) 0.7 30.0 (1.2) 2.4 100.0 (4.0) 6.6 300.0 (12.0) 24.0

mixtures, and we believe that many nitro products observed in smog chambers may reflect heterogeneous

reactions either during the actual chamber

reaction or during trapping out of the products. Since phenolic compounds are especially susceptible

to heterogeneous nitration, the origin of nitrophenols must be interpreted with extreme caution.

4. Reactions of Initial Products Benzaldehyde Reactions . Two processes appear

to be important for the reaction of the benzal- dehyde formed from toluene in the atmosphere: the reaction with OH and photolysis.

Niki et al . [16] recently reported the rate constant for the reaction

as k = 1.3 x 10" 11 cm

3 molec

_1

s _1

, which is identical within the experimental uncertainties to rate constant for other aldehyde-OH reactions. Addition of OH to the ring as observed for toluene

(reaction (2))is expected to occur no faster than addition to benzene, where kgn

1 c 1 T ' '- -

J-

' ' • - 1.2 x 10" 12 cm 3 molec J s

1

. Thus the attack of OH is expected to be largely at the aldehydic position.

Using our discharge flow systems, we have found that the reaction produces phenol as the only gas phase product [17]. Thus the initial reactions of

the benzaldehyde with OH is

cut) c:i) p —

followed by the reactions

CO C(O)0

2

- C(0)0- (t5)-2*-{0) —[61

p]—ro roj -"- pi

.

87
In our flow system phenoxy is converted to phenol,

presumably by a wall reaction with benzaldehyde

[17]. Considerable amounts of wall products are also found, but field ionization mass spectrometry

indicates that the products are largely various

states of oxidized phenol including as much as

10 percent nitrophenol.

Niki et al . [18] found o- and ^-nitrophenol in the Cl 2 -catalyzed reaction of benzaldehyde in presence of air and N0 2 . The benzoyl radical is first formed by the reactions of CI atom

I'hcHo + ci- un

(12) which is followed by reactions (9) and (10). A significant fraction of the benzaldehyde appears

as nitrophenols, presumably by reaction of phenoxy radical with N0 2 .

or . ro

The second important reaction of benzaldehyde,

photolysis, has been studied by Berger et al . [19]

in the gas phase and in the absence of oxygen. Three photolytic reactions are possible.

I'hH + CO

I'll- + HCO

(15)

(16)

(17) Although reaction (15) is energetically favorable

at all wavelengths of the visible spectrum, its measured quantum yield is significant only at wave- lengths less than 300 nm [19]. Reaction (16) is energetically possible only below 300 nm. Thus,

only reaction (17), which has an energy cut-off

at 330 nm, appears to be important in the solar

spectrum. However, the possibility of the generation of a triplet excited state that reacts with oxygen above 330 nm cannot be ruled out.

Such a reaction would be an important source of radicals and therefore critical for modeling

purposes

.

In both the reaction paths for benzaldehyde,

the phenoxy radical is eventually formed. In the atmosphere, the fate of phenoxy, we believe, is determined by the reaction with oxygen as well as by reaction (13).

However, thermochemical calculations indicate that

the DH°(C-0 2 ) in the resulting peroxy radical is small and that the reaction will be reversible.

The importance of reaction (19) therefore depends

on how fast the peroxy radical is trapped bv reaction with NO.

(20)

The competition between the formation of nitrophenol (reaction 13) and reaction (20) is

k n[PhO-][N0

2]

2.1

['' ][N°]

19 .

it ] [no] The following estimates are applicable: k 13 = 1.7 x 10 12 cm

3 molec" 1 s" 1

, k 20 = 5.0 x 10" 12 cm molec 1 s \ k 19 = 6.0 x 10~ 13 cm 3 molec -1 s" 1 and [0 2 ] = 0.01 M in air at 1 atm. If k_ 19

- 10 s which is consistent with our estimate of DH°(C-0 2 ) = 10 kcal/mol, the two processes will compete equally at N0 2 /N0 =1. At very high ratios of N0 2 /N0, however, the formation of nitrophenol will predominate. If the DH°(C-02 ) is much weaker than 10 kcal/mol (k 19 » 10 7 s"

1 ), reactio

(13) will predominate under most atmospheric

conditions, but if it is much stronger than

10 kcal/mol (k 19 » 10 7 s" 1 ), reaction (13) will

be unimportant and reactions (19) and (20) will predominate. The relative importance of reactions

(13) and ( 1 9)- (20) is very critical because

reaction (13) is a termination reaction whereas

( 1 9

) - (20) will lead to ring degradation and further oxidation of NO by reactions of the following type.

read

route

CO, + HCXHCHn

(23]

This reaction sequence is speculative, although

each step can be justified in most cases by

analysis of competing reactions. It does suggest

that a-dicarbonyl compounds should be important

secondary products. These compounds absorb light very strongly in the solar spectrum and can be significant source of radicals [20,21].

Cresol Reactions . The reaction of OH plus o-cresol was studied by Perry et al . [22] over the temperature range 300 to 435 K (reactions for £-

and m-cresols are expected to be similar). Rate constants were reported for two processes: (1) a nonreversible reaction believed to be hydrogen abstraction (k = 2.6 x 10" 12 cm

3 molec" 1 s" 1

) and

(2) a reversible reaction believed to be addition

to the ring (k = 3.1 x 10" 11 cm

3 molec" 1 s"

1 ). We postulate the following reaction pathways.

1( wet
I

Reaction (26) should lead to hydroxybenzaldehyde,

reaction (27) should be followed by reactions of

the type proposed for the simple phenoxy radical,

and reaction (28) will lead to dihydroxytoluenes.

(a)

Fig. 1. (a) Simulation of SAPRC EC-86 Toluene (* = experimental, T simulation)

.

We are beginning to investigate the products of reaction of o_-cresol plus OH in our discharge

flow system to determine if the above reaction

routes are valid.

Other homoegeneous reactions of cresols to be considered are the reactions with 0 3 and 0 atoms.

For ja-cresol we have obtained a second-order rate constant for reaction with ozone equal to about

1.4 x 10" 18 cm 3 molec" 1 s" 1

. At 0.05 ppm 0 3 , this reaction is about 1 percent of the OH-cresol

reaction, assuming the o- and p_-cresol have the

same reactivities. While the reaction may prove unimportant as a loss mechanism for cresols, it can be a dominant source of free radicals at high ozone concentrations if it produces radicals efficienctly . We hope to determine if this is the case in our studies of cresol-0 3 reactions.

Atkinson and Pitts [23] studied the reaction of

0 atom plus o_-cresol and found it to have a rate constant of 5.8 x 10" 13 cm

3 molec" 1 s"

1

. Since

the OH reaction is 100 times faster and since OH

is 100 times more abundant than 0 atom, this reaction with o-cresol is insignificant.

5. Modeling of Toluene Smog Chamber Data

The Statewide Air Pollution Research Center

(SAPRC) at the University of California, Riverside

has carried out a series of runs with toluene in their smog chamber facility. Concentrations of

toluene range from 0.2 to 2.0 ppm while the NO concentration was varied from 0.1 to 1.0 ppm.

xWe

have developed a mechanism to simulate these data. The mechanism includes the standard inorganic

reactions and those organic reactions which have

been discussed in previous sections. We have also

included reactions for formation and decomposition

of the major peroxynitrates as well as the termination reactions of H0 2

' with R0« and R0 2

* radicals.

Figures 1 and 2 show simulation and experimental

data for SAPRC Runs EC-77 and EC-86. Run EC-77

was with 0.28 ppm toluene and 0.58 ppm NO , which

run EC-86 was with 1.09 ppm toluene, 0.49x ppm NO and 0.16 ppm formaldehyde. The agreement betweeft the simulation and experiment data is very good.

(b)

cc c cc c

0 SO 100 ISO ZOO . 250 300 350 ftOO Time (min)

Fig. 1. (b) Simulation of SAPRC EC-86:

NO (* = experimental, I = simulation) and

NO2 (+ = experimental, 2 = simulation).

(c)

33 11)33 3 3 3 3 50 100 ISO ?00 ?50 300 350 Time (min)

Fig. 1. (c) Simulation of SAPKC EC-77: Ozone (* = experimental, 3 = simulation)

and Formaldehyde (+ = experimental,

F = simulation)

.

89
(d) (c)

350 (.00

Fig. 1. (d) Simulation of SAPRC EC-77: Fig. 2. (c) Simulation of SAPRC EC-86: Cresol-total (C = simulation), Dihydroxy Ozone (* = experimental, 3 = simulation) toluene (D = simulation, and Benzaldehyde and Formaldehyde (+ = experimental,

(B = simulation). F = simulation).

(a)

Time (min)

Fig. 2. (a) Simulation of SAPRC EC-86: Toluene (* = experimental, T = simulation).

(b)

Time (min)

Fig. 2. (b) Simulation of SAPRC EC-86:

NO (* = experimental, 1 = simulation) and

N0 2 (+ = experimental, 2 = simulation).

(d)

8 hohrhh ea 8B8BHB8B

Fig. 2. (d) Simulation of SAPRC EC-86: Benzaldehyde (+ = experimental, B = simulation) and PAN (* = experimental, P = sim- ulation )

.

There are discrepancies in the formaldehyde and

PAN values which in part may be due to experimental uncertainties. The reason for the over prediction

in EC-86 of the ozone near its maximum is not

clear. However, the effect is' also seen in the simulation of smog chamber data for other hydrocarbons, and thus may or may not be due to the

actual toluene mechanism.

6. Conclusions

The reaction of toluene in the atmosphere is very complex. We have a good understanding of the

initial reactions, but we still need to determine

the fate of the primary and secondary products.

We can model the toluene smog chamber reasonably

well, but out mechanism includes speculation

regarding many of the intermediates.

More smog chamber studies are needed to identify

the products of toluene reaction and their yields.

90
I Currently, data show very low material balances

which may be indicative of the formation of

aerosols or deposition of products on the chamber

'walls. The chamber data should be obtained over

a wide range of conditions because the sensitivity

of individual reactions varies with the conditions.

Thus, by using a wide range of conditions,

different parts of the model can be tested.

The current study of individual reactions of various intermediates should be continued. This

i work has been one of the most helpful sources of

information in developing the toluene mechanism.

Finally, the inability to simulate the ozone

data in smog chamber runs, indicate a need for a better understanding of the chemistry that controls

the ozone concentration. Since this effect appears

to be common to the simulation of data for other 'hydrocarbons, the problem may not be solely with

the organic part of the mechanism.

I References

[I] Calvert, J. 6., Environ. Sci. Tech . 10, 256

(1976).

[2] Lonneman, W. A., Kopczynski, S. L., Danley,

P. E., and Sutterfield, F. D., Environ. Sci - Tech . 8, 229 (1974).

] [3] Crabtree, J. H. , private communication.

[4] Heuss, J. M. and Glasson, W. A., Environ . Sci. Tech . 2, 1109 (1968).

[5] Davis, D. D., Bollinger, W. , and Fischer, S.,

J. Phys. Chem . 79, 293 (1975); Davis, D. D., Investigation of Important Hydroxyl Radical Reactions in the Perturbed Troposphere,

EPA-600/3-77-11 (October 1977).

[6] Hansen, D. A., Atkinson, R. and Pitts,

J. N . , Jr., J. Phys. Chem . 79, 1763 (1975).

[7] Perry, R. A., Atkinson, R. and Pitts,

J. N . , Jr., J. Phys. Chem . 81_, 296 (1977).

[8] Atkinson, R. and Pitts, Jr., J. N. , J. Phys . Chem . 79, 295 (1975).

i [9] Nakagawa, T. w., Andrews, L. J., and Keefer,

R. M. , J. Amer. Chem. Soc . 82, 269 (1960).

[10] Hendry, D. G. , Mill, T. , Piszkiewicz, L.

,

Howard, J. A., and Eigenmann, H. K., J. Phys . Chem. Ref. Data 3, 937 (1974).

[II] Doyle, G. J., Lloyd, A. C, Darnall, K. R. , Winer, A. M. , and Pitts, J. N. , Jr., Environ . Sci. Tech . 9, 237 (1975).

I [12] Kenley, R. A., Davenport, J. E., and Hendry,

D. G. , J. Phys. Chem . 82, 1095-1096 (1978).

[13] Kenley, R. A. and Hendry, D. G., manuscript

in preparation.

ft

[14] O'Brien, R. J., Green, P. J., and Doty, R. A., Interaction of Oxides of Nitrogen with Aromatic Hydrocarbons, 175th National Meeting, of the American Chemical Society,

March 1978.

[15] Fitz, D. R., Grosjean, D. , Van Cauwenberghe,

K. , and Pitts, J. N. , Jr., Photo-oxidation

Products of Toluene-NO Mixtures Under

Simulated Atmospheric Conditions, 175th Meeting of the American Chemical Society, March 1978.

[16] Niki, H. , Maker, P. D., Savage, C. M. , and Breitenbach, L. P., J. Phys. Chem . 82, 132 (1978).

[17] Kenley, R. A., Lan, B., and Herdry, D. G., unpublished data.

[18] Niki, H. , Maker, P. F., Savage, C. M., and Breitenbach, L. P., Fourier Transform IR Studies of Gaseous and Particulate Nitro- geneous Compounds of Atmospheric Interest, 175th National Meeting of the American

Chemical Society, March 1978.

[19] Berger, M. , Goldblatt, I. L., and Steel, C,

J. Amer. Chem. Soc . 95, 1717 (1973).

[20] Porter, G. B. , J. Chem. Phys . 32, 1587 (1960).

[21] Bouchy, M. and Andre, J. C, Molec. Photochem . 8, 345 (1977).

[22] Perry, R. A., Atkinson, R., and Pitts, J. N., Jr., J. Phys. Chem . 81_, 1607 (1 977).

[23] Atkinson, R. and Pitts, J. N., Jr., J. Phys . Chem . 79, 541 (1975).

Summary of Session

The presentation by Hendry emphasized the importance of aromatic compounds in the chemistry

of urban air pollution. Single ring aromatic

compounds account for 25-40 percent of the carbon species found in urban air. From chamber studies

these comoounds are known to be reactive in the production of ozone (0 3 ). Therefore a knowledge

of the atmospheric chemistry of simple aromatics

is required for inclusion of these compounds in tropospheric models to predict their role and contribution to photochemical smog formation. The importance of a better understanding of the chemistry was illustrated in comments by Atkinson

and Hendry on the uniqueness of the 0 3 formation curve and the current inability to simulate 0 3 smog chamber data.

The major theme of the discussion and the majority of the uncertainties centered around mechanisms of reactions of primary and secondary

aromatics in the atmosphere. There was general

agreement that the initial reaction can be accounted for almost solely by attack of the hydroxyl (OH) radical. For methyl substituted benzenes, the accepted mechanisms are hydrogen abstraction at the methyl group and OH addition at the ortho

position. However, there was a degree of

91
uncertainty regarding the branching ratio for the two pathways, with the only significant work being done on toluene. In addition Atkinson pointed out

the thermodynamic favorability of OH addition at the methyl position. Inclusion of this pathway

could alter mechanistic consideration of other

i nvestigators.

There was considerable discussion and some lack of agreement on product yields. For toluene, O'Brien reported lower product yields than Hendry for benzaldehyde and o-cresol by factors of approximately 4 and 10 respectively. There was some speculation that Hendry's results were higher because he based his yields on the total amount of gas-phase carbon analyzed. In any event none of

the investigators have yet analyzed aerosol carbon

or carbon on the walls of reaction chambers. Both O'Brien and Atkinson reported carbon balances well below 100 percent.

The question of the mechanism(s) of ring opening

was raised several times. Tentative mechanisms

were proposed in the papers by Hendry and Atkinson. Evidence for ring opening was given by the observation of peroxyacetyl nitrate and carbon monoxide by Atkinson. Ring opening could be of considerable importance as a source of free radicals and simple oxgenated products.

The status of the uncertainty on reaction mechanisms may be illustrated by the fact that the only work reported on basic mechanisms was the low pressure flow tube studies of Hendry. Most

of the mechanistic work reported was on toluene. The reactions and fate of aromatic products was largely unconsidered. In his paper Hendry

discussed the attack of OH on benzaldehyde and the cresols and subsequent reaction pathways. With regard to reaction kinetics, relatively

good agreement was reported for the OH reaction with the primary aromatics-benzene , toluene and xylenes. Little data are available on the higher

homologues or on OH reaction with aromatic

products. O'Brien reported ratio measurements,

relative to toluene, for o-cresol and benzaldehyde. Atkinson presented reaction rate data for the cresols.

It was evident from the discussion that some problems exist with regard to analytical measure- ments of products. All of the analyses reported

during the discussions were performed by gas chromatography (GC). O'Brien reported some difficulty with some product measurements at low concentrations, e.g. cresols. No analyses were

reported by other techniques such as mass spectroscopy or Fourier transform infrared

spectroscopy. Either of these techniques could give better time resolution and the possibility

of observing intermediates. Finally there is the larger question of the amount and nature of products in the aerosol phase.

Comments

Roger Atkinson, Statewide Air Pollution Research

Center, University of California, Riverside, California 92521

I would like to make three points:

1) Besides the two initial reactions of the o( radical with the substituted aromatics (taking toluene as an example),

(ID

OH radical addition at the 1-position leads to' the formation of phenol and CH3 radicals

(in)

h

Elimination of CH 3 from radical III can be calculated to be ^9 kcal mol

-1 endothermic. This together with an activation energy for the addition of CH 3 radicals to toluene of ^4 kcal mol

-1 [1], leads to an activation energy of ^ 13 kcal mol

-1 for reaction (3). Hence reaction [j (3) will be favored over elimination of an OH radical (analogous to reaction (1)) from this OH-toluene adduct. The occurrence of this reaction pathway would hence mean that the values 1, of ki and k 1/ ( k 1 + k 2 ) obtained by Perry, Atkinsoif and Pitts [2,3] are upper limits. This may be especially true for o-xylene where, by analogy

with the 0(

3 P) atom reaction [4], OH radical addition at the methyl substituted positions is likely to be appreciable, and for which the repor

ed value of k 1/ ( k 1 + k 2 ) appears to be high, with

a low value of Ei e , compared to the other

aromatic hydrocarbons.

2) At the Statewide Air Pollution Research

Center, University of California, Riverside, we

[5] have recently determined rate constants for

the reaction of OH radicals with 0-, m- and p-cresol from the rates of disappearance of the cresols and n-butane in irradiated NO -organic-ai mixtures of atmospheric pressure and 3oo ± 1 K. Using a value of k(0H + n-butane) of 2.73 x 10" 12 cm

3 molec 1 s"

1 at 300 K [6] rate constants k (cm 3 molec

-1 s" 1 } of (4.7 ± 0.4) x 10~ 12

; (6.7 ± 0.7) x 10" 12 and (5.2 ± 0.5) x 10" 12 were obtained [5] for o-cresol, m-cresol and p-cresol. Further experiments [7] have shown that the NO photooxidations of the cresols form hydroxyx nitroluenes as the major observed gas phase

92
iaromatic products, the 2,3- and 2,5-isomers from p-cresol, the 3,2- and 3,4-isomers from m-cresol

,

land the 4,3-isomer from p-cresol.

3) Ring cleavage in the OH-aromatic systems

may occur, in part, through the reaction sequence

(ii)

Robert J. O'Brien, Patrick J. Green, and Richard

M. Doty, Department of Chemistry, Portland State University, Portland, Oregon 97207

The product formation and dynamics for the toluene (T) system under ambient conditions are of great current interest. For the simple case where

hydroxyl radical determines both formation and

loss of "stable" products (P.):

T + OH a

i Pi

P. + OH loss

(III) Radical (III) may then react with NO to form N0 2 and

the radical (j^J^

H wnicn would probably undergo

ring opening, leading to a variety of oxygenated

species. H atom abstraction from radical (II) by

0 2 to form o-cresol is ^ 26 kcal mol" 1 exothermic,

with the bond strength of the C-H bond at which

labstraction takes place being °» 20 kcal mol" 1

.

From group additivity calculations, using a C00-H bond energy of 90 kcal mol

-1 (the same as that for H00-H, AH- (III) <\> -6 kcal mol" 1

. As My II) ^ 1 kcal mil" 1

, formation of (III) from

(II) is ^ 7 kcal mol" 1 exothermic and hence

radicals (III) and (II) will be in equilibrium.

Hence this reaction pathway leading to ring

opening is expected to become more important at lower temperature, and vice-versa.

References

[1] Cher, M. , Holl ingsworth, C. S., and Sicilio,

F., J. Phys. Chem. , 70, 877 (1966).

[2] Perry, R. A., Atkinson, R. , and Pitts, J. N.

,

Jr., J. Phys. Chem. , 81_, 296 (1978).

j[3] Perry, R. A., Atkinson, R. , and Pitts, J. N.

,

Jr., J. Phys. Chem. , 81_, 1607 (1978).

[4] Grovenstein, E. , Jr. and Mosher, A. J.,

J. Amer. Chem. Soc , 92, 3810 (1970).

[5] Atkinson, R., Darnall , K. R. , and Pitts,

J. N., Jr., J. Phys . Chem. , submitted for publication (1978)!

[6] Perry, R. A., Atkinson, R. , and Pitts, J. N.

,

Jr., J. Chem. Phys. , 64, 5314 (1976).

[7] Darnall, K. R., Atkinson, R. , Glangetas, A., Winer, A. M. , and Pitts, J. N. , Jr., unpublished data (1978).

the maximum production concentration will be given

by

(1)

[ p

max-j a.i_

form N0 2 and

[T] R

i where a. is the yield of product P. in the primary

reaction step and R. = b./a is the ratio of the OH rate constant with the product, P., to the rate constant with toluene.

If we assume pseudo first order loss for toluene (constant OH concentration) we may

integrate the rate expression to determine the length of time required to achieve maximum

concentration (tTax

) in terms of the toluene

lifetime (t t

). this is given by

^max

i

T

(RrD

In R, (2)

A plot of this function is given in figure 1 100

1

Fig. 1. Variation of the time for a product to reach maximum concentration relative to the toluene lifetime (t

t!iax /T

T

) with the relative reactivity with OH radical (R.)

as given by eq. (2).

1

93
From this plot we may determine (for example)

that for a typical atmospheric lifetime of toluene of 10 hours the concentration of p-cresol

(which reacts six times faster than toluene with

OH) should reach a maximum in 3.5 hours. If the yield of p^-cresol is 5 percent (see below) and

the ambient toluene concentration is .020 ppm, this maximum concentration calculated from eq. (1)

is about 0.2 ppb.

Assuming pseudo first order toluene loss the variation of a product with toluene concentration

is given by

R-l

R-l

(3)

The general form of this equation for various

values of is shown in figure 2.

R =0.1

Toluene ppm

Fig. 2. Variation of a product concentration with

toluene concentration for different values of R.j , assuming a = 10 percent.

The values of a. may be obtained from the individual rate measurements of Atkinson and others. However, considering the errors present

in each measurement a separate measurement of the ratio itself may be preferable. We have made

such measurements by irradiating a mixture of P. and toluene at about a 10 to 1 ratio (P./T). A

1 plot of In P. vs. In T gives the value of R . as the slope. This analysis is not sufficient^

the product photolyzes to any appreciable extent.

For the case of o-cresol the photolytic lifetime

in our reaction vessel is 10" minutes and is probably negligible.

For the more general case of non first order

toluene loss (variable OH concentration) we may

still derive an expression to analyze production

formation and loss. For the same mechanism given

above (reactions a and b) it can be shown that

A[P.] + R. [T] d[T] «

i A[T] (4)

For the simple case where a product is totally

unreactive, R. = 0 and a plot of P. vs. T will give a straight line with slope = a. For the case where the product does react further, the second term on the left hand side of eq. (4) corrects for this loss of product. The variation

of ^-cresol and of benzaldehyde for one of our experiments are shown in figures 3 and 4. This experiment was carried out by irradiating toluene

and N02 each at about 4 ppm in a 250 L evacuable

glass vessel with a mixture of fluorescent black

lights and sun lamps.

The yield of each product may be determined

from the slope of these plots. For benzaldehyde

we obtain a 2.5 percent yield and for o-cresol a

5 percent yield. These yields are much lower

than those measured by Hendry in his low pressure

flow system.

We have been initially skeptical of our low yields, especially for o-cresol since it is about

ten times lower than the yield reported by Hendry.

To double check this result we have carried out experiments in which we start with a mixture of toluene and p^-cresol (4 ppm and 1 ppm) respectively). The decay of p^-cresol is then modified by formation of p^-cresol from toluene. Equation (3) holds for any initial product concentration so we I have plotted the data for this experiment in the required form in figure 5. The yield of jg-cresol

j is found from the slope to be 5 percent, in agreement with the other experiments. This experiment has the advantage of generating a large!

160 r—

n

a.

Q.O

in

91

4000 500 Toluene ppb

Fig. 3. Analysis of cresol formation from toluene. Lower data points are a plot of o-cresol

vs. toluene. Upper data points with line?

are a plot of (C) + 6 / (C)/(T) d(T) vs.

(T) in accordance with eq. (s). (T = toluene; C = p^-cresol).

94
60i—

Slope = .025 n

Q.

a.

1000 2000 3000 Toluene ppb

4000 5000

Fig. 4. Analysis of benzaldehyde formation from

toluene. Lower data points are a plot of benzaldehyde vs. toluene. Upper data points with straight line are a plot of

(B) + 2.3 / (B)/(T) d(T) in accordance

with eq. (3). (T = toluene; B = benz- aldehyde) .

3 i10 11 12 Toluene ppm

Fig. 5. Analysis of competitive disappearance of

a mixture of 11 ppm toluene + 2 ppm

o-cresol . Data plotted as (C) + 6 / (C)/ CO d(T) in accordance with eq. (3).

(T = toluene; C = o-cresol).

unambiguous, initial cresol peak with the gas chromatograph.

We currently have no explanation for the lower product yields for toluene but the most obvious explanation would lie in the difference in pressure between the two measurements. Our yields

of nitrotoluene are in agreement with those obtained from Hendry's work and would be negligible

at sub ppm N02 concentrations. Our mass balance

for gaseous toluene products is then well below

100 percent. The carbon balance may possibly be accounted for by aerosol analysis which we will do

in the future.

A. R. Ravishankara, Applied Science Laboratories,

EES, Georgia Institute of Technology, Atlanta,

Georgia 30332

We have noticed a decrease in the net rate constant

for m-xylene (drops at lower pressures e.g. ^ 3 Torr

of Ar). In Hendry's experiments there could be an overemphasis on the abstraction route since his flow

tube pressures are not very high.

Recommendations

Importance

Study of the atmospheric reaction processes of aromatic hydrocarbons is in its early stages. Our current knowledge about these compounds is rather primitive compared to the alkanes and alkenes. However, aromatics are major components of urban atmospheres and elucidation of their reaction

pathways is essential for the following reasons.

1) Oxidant formation . As major urban hydro- carbons which react relatively rapidly with

hydroxyl radical, aromatics will contribute

directly to ozone formation and buildup and they seem to generate appreciable quantities of PAN, itself a harmful oxidant.

2) Direct health effects . Oxidant products of aromatic hydrocarbons are poorly characterized

and present a potential health hazard of undetermined magnitude.

3) Aerosol formation . Gas phase mass balances

for smog chamber experiments with aromatics are very poor and may indicate appreciable aerosol formation. If so, the aerosol so formed may contribute to a heterogeneous component of tropospheric chemistry which is currently unrecognized;

this heterogeneous component may well impact other

areas - in particular N0

X conversion to nitric acid or free radical loss processes.

Current Status

1) Rate constants . Considerable work has been done to determine the reaction rate constants for hydroxyl radicals with the chief aromatic constituents of the atmosphere. Agreement between

various groups is quite good so this question is resolved. Ozone and other free radicals (HO2,

N0 3) are known to react slowly with aromatics and

are therefore, at present of minor importance. Relative rates of ring addition versus side chain abstraction, while less certain than the overall rates, are also fairly well settled. Rate constants for reaction of OH with some of

the more important reaction products of aromatic hydrocarbons (cresols, benzaldehyde, etc.) have

95
also been measured.

2) Product identification . Major products of

the reaction of OH with toluene which have been identified are the following: cresols, nitro- toluenes, benzaldehyde, benzyl alcohol, benzyl

nitrate, peroxybenzoyl nitrate, peroxyacetyl

nitrate, and carbon monoxide. Yields of these compounds have been determined at low pressure and' are becoming available at atmospheric pressure as well. Currently the low pressure yields are considerably higher on an absolute basis than those at high pressure, but on a relative basis are in good agreement. Products have been identified for the reactions of some other

aromatics as wel 1

.

Products of the subsequent reactions of the primary products are known in a few cases.

The hydroxyl radical is a key component in controlling loss rates of the primary products but other processes, such as direct photolysis or reaction with 0 3 , R0 2 , RO, N0 3 , etc. may be important as wel 1

.

3) Ozone formation . Only a limited amount of work has been done on modeling the ozone profiles

in aromatic/NO or mixed hydrocarbon/NO

x systems which include aYomatics, because of the general lack of knowledge about the detailed reactions

involved. However, the limited work done to date indicates that ozone profiles are different than those in nonaromatic systems and in some cases are difficult to model unless unique radical-radical

reactions are invoked.

4) Analytical techniques . Current studies of aromatic hydrocarbon systems are severely hampered

by a lack of versatile techniques for analyzing

the high molecular weight products involved. Techniques which have been employed include gas chromatography, gas chromatography-mass spectrometry and to a limited extent Fourier transform

infrared spectrometry. These techniques are difficult to employ when they are successful, and

are often unsuccessful . Much time has been spent

in adapting these techniques for use in the study

of aromatic hydrocarbons, but they still suffer from some inherent problems.

Recommendations

1) Absolute yields of the major known primary

products of aromatic-OH reactions should be determined at atmospheric pressure. These aromatics would include as a minimal set benzene,

toluene, the xylenes, trimethyl benzene and some

alky! benzenes such as ethyl benzene.

2) Rate constants for the various processes

these products undergo should be determined.

Although a large number of compounds are involved,

competitive kinetic studies employing several compounds simultaneously may suffice in some cases. This would reduce the total number of necessary

experiments.

| 3) A carbon mass balance for the gaseous products including CO and CO2 should be obtained

for the major aromatics. The mass balance should ! include the carbon content of any aerosol formed.

4) New analytical techniques should be investi- gated for application to the study of aromatics. These techniques would be doubly useful because

they would be equally applicable to the study of higher moleculas weight alkanes and alkenes. Techniques which might be investigated include improvement of gc sampling techniques and separation efficiency on the column, direct mass spectral

analysis employing non-framentation ionization,

liquid chromatography, and field desorption mass spectrometry.

In all these techniques every attempt should

be made to work at realistic reactant concentrations and total pressures and to induce minimal sample alteration. However, some low pressure techniques may have to be employed (e.g., direct

ms sampling) because of the lack of any other viable alternatives for direct analysis of inter- mediates. The current advancement of knowledge

in this area is now limited by analytical

methodology. Advancement of knowledge in the alkane and alkene systems will soon suffer the same fate, as the chemistry of the low molecular weight compounds becomes worked out, and higher

members of the series are studied.

5) Heterogeneous processes may be of great

importance in the aromatic hydrocarbon systems. The impact of these processes may well extend

beyond the purely aromatic systems and influence

the chemistry of NO and of free radicals generated from other classes of compounds. An attempt should be made to assess the significance

of these processes on the overall chemistry of the troposphere.

6) Compouter modeling of the aromatic hydro- carbon system should be continued in order to assess the ozone forming potential of these hydrocarbons. It is expected that these modeling

efforts will become more meaningful as more

basic rate and product data become available.

Recommendations 1 to 3 may be expected to be completed with current funding in the next year

or two. Recommendations 4 and 5 are much more

ambitious and will require a long term committment

and considerable additional funds for instrument

development.

96
Session VI

S0X Chemistry

'UU

SO

x CHEMISTRY

Jack G. Calvert

Chemistry Department

The Ohio State University

Columbus, Ohio 43210

An evaluation has been made of the existing kinetic data related to the elementary,

homogeneous reactions of S0 2 within the troposphere. A set of recommended values of the rate constants for these reactions is presented. The results show that the direct photooxidation

of S0 2 by way of the electronically excited states of S0 2 is relatively unimportant for most conditions which occur within the troposphere. The oxidation of S0 2 within the natural troposphere is expected to occur largely by way of reactions 39, 31, and 33, with reaction 39 being the dominant path: HO + S0 2 (+M) -> H0S02 (+M) (39); H0 2 + S0 2

-> HO + S0 3 (31); CH 3 0 2 + S0 2

->- CH 3 0 + S0 3 (33). For certain special conditions within the troposphere the oxidation

of S0 2 by way of the products of the ozone-olefin reaction may be significant. Also the reaction of 0( 3 P) with S0 2 may contribute somewhat to the S0 2 removal in N0 2 -rich, 0 2 -deficient

stack gases in sunlight during the early stages of dilution of the plume.

The complete paper upon which most of the talk was based is: "Mechanism of the Homogeneous

Oxidation of Sulfur Dioxide in the Troposphere", by Jack G. Calvert, Fu, Su, Jan W. Bottenheim,

and Otto P. Strausz which appeared in the Preprint Volume I, Plenary Papers, International Symposium on Sulfur in the Atmosphere, July, 1977, and which was presented at the Symposium on September 7-14, 1977, Dubrovnik, Yugoslavia. The paper has been published in Atmospheric

Environment , 1_2, 197 (1978).

Key words: Kinetics; photochemistry; review; sulfur dioxide; troposphere.

Summary of Session

The discussion was concerned with two basic

problems - what is the mechanism of conversion of

S0 2 in the atmosphere, and what do we know about

aerosol formation arising from SO and NO reactions. There appeared to be § consensus that

the reaction of S0 2 with OH is the most important

homogeneous mechanism, but there is still great

interest in quantifying the role of H02 , R0 2 , and

the Criegee intermediate in this process. There

is considerable experimental work underway on aerosol formation, the incorporation of NO in aerosols, and the role of specific radical^ in aerosol formation.

Whitten opened the discussion with a descriptions of his modeling results for the Los Angeles

basin. He pointed out that while OH levels

could be reduced by reducing hydrocarbon or NO

levels, they were relatively constant for a given

HC/NO ratio. Thus if the reaction of OH with

S0 2 controls the S0 2 level, control of the HC or

NO levels will not necessarily have any affect

on the OH levels if the HC/N0

X ratio remains

fixed. A similar conclusion was reached in the case of H02 radical levels, although the new

value for the rate of N0X + H0 2 means that H0 2 levels are reduced to the point where the

H0 2 + S0 2 reaction is probably unimportant.

Whitten also reported modeling studies which

suggest that a fairly rapid conversion of S0 2 to sulfate takes place in fog droplets. The observations used in this study were based on one days sulfate collections from 14 monitor stations (24 hour averages). A photochemical model was not sufficient to account for the S02 conversion rate.

Miller reported on smog chamber results which supported Whittens observations. He also discussed measurements which show that the rate constant for

H02 + S0 2 is no greater than 1 ppm

-1 min -1 and for

CH 3 0 2 + S0 2 is no greater than 2 ppm

-1 min The

role of N0 X in aerosol formation was also discussed.

Heicklen suggested that excited molecules such

as N0 2 , formed by irradiation at wavelengths above

400 nm, might react with S0 2 . Also photo excited aldehydes and ketones might be important. Cox reported on the photolysis of HONO in the presence

of S0 2 . The end product was an aerosol. With

99
added NO some of the NO was incorporated in the

aerosol , x but since NO Cctn lead to aerosol formation without S0 2 being present, it was difficult

to say how much of the lost N0 X was incorporated.

Jeffries reported smog chamber results on the effect of added CO on SO2 conversion rates in natural background air, which indicate that OH is the only important oxidizing species for S0 2

. Ravishankara noted that aerosols are readily

formed in uv flashed H 2 0-S02 mixtures, probably

by photolysis of H2O. Huie raised the question

of the reactivity of the Criegee intermediate with

respect to S0 2 , pointing out that if the Criegee intermediate isomerized to a dioxirane intermediate

it probably decomposes through a "hot" acid or ester before it has time to react with S0 2 . However, Calvert noted, one would expect for larger olefins an increasing chance of stabilizing

the Criegee intermediate and observing some direct

chemistry.

Niki noted that generating HO2 by reacting CI with H 2 in the presence of 0 2 and S0 2 led to H 2 0 2 but no conversion of the S02- If NO was added

(to drive the H0 2 to OH), then sulfuric acid

aerosol is formed. In the case of CH 3 0 2 some conversion of S0 2 is observed and there is evidence to suggest that sulfones are possibly

formed. Jeffries asked if Niki has any evidence

that water intercepted the Criegee intermediate to make acetic acid. Niki has not yet carried out

an experiment with added H 2 0. Jeffries noted that

in earlier work they had not seen acetic but had

seen formic acid in the reaction of ozone with propylene.

Comments

David F. Miller, Battel 1 e-Col umbus Laboratories,

Columbus, Ohio 43201

Gary Whitten presented modeling results relating

OH concentrations to initial concentrations of NMHC (nonmethane hydrocarbons) and NO . According

to his model, OH concentrations are predicted to be nearly constant for any NMHC/N0

x ratio. I'd like to add that our smog chemaber results,

reposted last year in Dubrovnik, led to the same conclusion. This finding has a very important

implication regarding precursor controls designed

for limiting ozone. Because S02 competes with

NMHC and NO for OH, proportional control of NMHC

and NO cou^d result in an increase in the conversion of S0 2 to sulfate.

Secondly, I'd like to comment on some smog chamber experiments in which we've irradiated mixtures of nitrous acid with S0 2 and either CO ur

CHl, to estimate the S0 2 oxidation rates attributable to H0 2 and CH 3 0 2 . Although our analyses of

the data are less than satisfactory, primarily

because of so much uncertainty about the nitrogen

oxides chemistry, we estimate that the rate for

H0 2 + S0 2 is not greater than 1 ppm' 1 min" 1 and the

rate for CH 3 0 2 + S0 2 is not greater than 2 ppm" 1 min 1

When we irradiate just S0 2 with nitrous acid we observe an NO loss in excess of that for the experiment without S0 2 . The amount of the NO loss corresponds to the amount of H2S0,, formed, w£ suspected that either NO or N0 2 might be incorporate

ed in the aerosol phase. It has been suggested,

for example, that HS0i« might react rapidly with

N0 2 to give aerosol mixtures of sulfuric and nitric

I

acids. However recent chemical analyses of filter : collections from such reactions show very low nitrate levels relative to sulfate. Thus, if such 1 reactions occur, the nitric acid apparently ends up!

in the gas phase.

I The experiments were conducted with the followinc concentration ranges: S0 2 (400-500 ppb), HN0 2 (100-'; 200 ppb), NO and N0 2 (20-200 ppb), CO (2-200 ppm) and CHl, (200 to 900 ppm).

Richard A. Cox, U.K.A.E.A., Environmental and Medical Sciences, A.E.R.E, Harwell, Oxfordshire

0X12 0RA, England

As part of an extensive research program on kinetics of reactions of S0 2 in the atmosphere,

we have used the photolysis of H0N0 to produce

OH radicals and allowed them to react with S0 2 at 25 °C and 1 atm pressure in air. We have drawn the following conclusions:

(a) Reaction of OH with S0 2 occurs with a rate constant of 'v 6 x 10~ 13 cm

3 s _I

(b) This reaction proceeds by addition to give H0S0 2 and the subsequent radical chemistry

leads to a short chain reaction in which NO is oxidized to N0 2 . This chain reaction is \ inhibited by N0 2 . 5

(c) The final reaction product is an aerosol,

physically resembling model H 2 S0i* aerosols.

(d) In the N0 2 - inhibited system, the aerosol contains SO -NO species yielding

equimolar proportions of S0i»

2 ~ and N0 3 ~ on hydrolysis in water.

Harvey E. Jeffries, Department of Environmental Science and Engineering, University of North Carolina, Chapel Hill, North Carolina 27514

In a series of experiments performed in UNC's outdoor aerosol chamber, S0 2 (at ^0.3 ppm) oxidation in natural background air (< 20 ppb

NC , < 50 ppb C organics) was followed by observing aerosol number (by CN), aerosol volume (by EAA) , and aerosol sulfur content (by XRF analysis

of filters). Runs were repeated with various additional amounts of CO added (5, 10, 15, 25 ppm). The additional CO resulted in delays in time of CN peak, small increases in 0 3 produced,

and reduced aerosol volume (by both EAA + XRF).

100
In one run, no CO was added initially, but when

i a steady rate of increase in aerosol volume had been established, 25 ppm of CO was injected;

aerosol volume production (i.e. growth) was

' totally stopped within 4 minutes. CO's role in this otherwise low concentration system is to convert OH to H0 2 . It seems clear that OH was

by far the major oxidizing species. It is expected that under urban conditions, however,

(i.e. higher NO concentrations) the effects of

I CO would not be observed because the higher NO converts H0 2 to OH.

A. R. Ravishankara, Applied Science Laboratories,

EES, Georgia Institute of Technology, Atlanta,

Georgia 30332

We have noticed formation of aerosols directly

in our system when ^ 300 mTorr of H 2 0 and S0 2 are

| photolyzed. This mechanism could be important where water concentrations are high i.e., very quick oxidation of sulfur dioxide leading directly

I to aerosol. (Even a mixture of S0 2 , 0 3 and H 2 0 gives aerosols). The water concentration needed

to get this aerosol formation seemed rather

magical - aerosols formed only after a critical amount of water was present.

Recommendations

It is recommended that kinetic and chemical data regarding S0 2 chemistry in the troposphere

be obtained. The classes of reactions are of six types, with the first four of these being almost equally important.

1. Of most importance is obtaining both product and rate information of HO S0 2 and R0 S0 2 with H 2 0, NO , 0 2 , hydrocarbons, NH 3 , and comoinations of these gases.

2. Of essentially equal importance is obtaining

information on the fate of S0 2 in 0 3 -olefin-0 2 reactions. There are three subsections of this

problem which should be attacked in the following

order:

a) characterize the intermediates which

which react with S0 2 b) obtain products and rate coeff ici.ents for the reactions of these intermediates

with S0 2 c) study the effect of adducts such as

H 2 0, NO , hydrocarbons, NH 3 , and combinations of these gases.

3. More data is needed on the rate coefficients

(and products) for the reactions of H0 2 , HO, and

0(

3 P) with S02 . These data should include pressure, temperature, and humidity studies. In the case of H02 , there is a large uncertainty in the rate constant. With regard to HO and 0(

3 P), fairly reliable values exist. However because of

the importance of the HO radical, which appears to be the most important species for SO2 removal, it

is important to have as accurate a rate coefficient

as possible.

4. The reactions of R0 2 radicals with S0 2 should be investigated to determine products and rate coefficients at a variety of pressures, temperatures, and humidities, and in the presence

of NO , 0 2 , NH 3 and hydrocarbons. The reactions

of R0 radicals with S0 2 appear to be unimportant

in the troposphere, and we do not give a high

priority to their study. However it would be useful to actually have rate coefficients for R0 reactions with S0 2 to know exactly what role these reactions do play.

5. A low priority recommendation is the study

of the possible reaction of electronically

excited N0 2 with S02 . There is no evidence that

a reaction occurs, but this should be confirmed.

6. The direct photoexcitation of S0 2 is not important in the removal of S0 2 in the troposphere,

and we do not recommend studies in this area. However we do point out that such reactions may be important in the formation of sulfur-containing

organic aerosols. If so then such reactions could

be of significance in aerosol chemistry.

101

Workshop Attendees

Larry G. Anderson

General Motors Research Labs. Environmental Science Department

Warren, MI 48090

Paul Ascher Northrop Services, Inc.

P. 0. Box 12313 Research Triangle Park, NC 27709

Roger Atkinson

Statewide Air Pollution Research Center University of California

Riverside, CA 92521

N. Basco University of British Columbia Chemistry Department Vancouver V6T 1W5 British Columbia, Canada

L. Batt Chemistry Department

University of Aberdeen

Meston Wal

k

Aberdeen, Scotland AB9 2UE

Sidney W. Benson Chemistry Department

University of Southern California

Los Angeles, CA 90007

Jack G. Calvert

The Ohio State University

Chemistry Department

140 W. 18th Avenue

Columbus, OH 43210

William P. L. Carter Statewide Air Pollution Research Center University of California

Riverside, CA 92521

C. Richard Cothern

EPA

1705 N. Stafford Street Arlington, VA 22207

Richard A. Cox U.K.A.E.A. Environmental and Medical Sciences

Division, A.E.R.E. Harwell, Oxfordshire 0X12 0RA England

R. J. Cvetanovic

National Research Council of Canada Division of Chemistry

Ottawa, Ontario, Canada K1A 0R6

Kenneth L. Demerjian

Environmental Protection Agency

NCHS-C Room 320-B

Davis Drive Research Triangle Park, NC 27711

William B. DeMore Jet Propulsion Laboratory

4800 Oak Grove Drive Pasadena, CA 91103

Marcia C. Dodge Environmental Protection Agency

MD-84 Research Triangle Park, NC 27711

William H. Duewer Lawrence Livermore Laboratory (L-262)

P. 0. Box 808 Livermore, CA 94550

A. Fontijn Aerochem Research Lab., Inc.

P. 0. Box 12 Princeton, NJ 08540

David Garvin

National Bureau of Standards

Washington, DC 20234

Lewis H. Gevantman

National Bureau of Standards

Office of Standard Reference Data Washington, DC 20234

David Mark Golden

SRI International

333 Ravenswood Drive Menlo Park, CA 94025

T. E. Graedel

Bell Laboratories

Room ID- 34

9

Murray Hill , NJ 07974

David Gutman Illinois Institute of Technology

Department of Chemistry

Chicago, IL 60616

Robert Hampson

National Bureau of Standards

A145, 222 Washington, DC 20234

Julian Heicklen

The Pennsylvania State University

Department of Chemistry

152 Davey Lab. University Park, PA 16802

103
Dale G. Hendry

SRI International

333 Ravenswood Drive

Menlo Park, CA 94025

John T. Herron

National Bureau of Standards

A145, 222 Washington, DC 20234

Jimmie A. Hodgeson

Office of Environmental Measurements

National Bureau of Standards Washington, DC 20234

Frank P. Hudson Department of Energy

(E-201) Washington, DC 20545

Robert E. Huie

National Bureau of Standards

A145, 222 Washington, DC 20234

Harvey E. Jeffries University of North Carolina Department of Environmental Sciences

and Engineering

Chapel Hill, NC 27514

J. A. Kerr Department of Chemistry

The University

Birmingham, B15 2TT England

Michael J. Kurylo

National Bureau of Standards

A145, 222 Washington, DC 20234

Stuart Z. Levine Brookhaven National Laboratory

51 Bell Avenue - Building 426

Upton, Long Island, NY 11973

Alan C. Lloyd Environmental Research and Technical Institute

2030 Alameda Padre Serra

Santa Barbara, CA 93103

Richard I. Martinez

National Bureau of Standards

A145, 222 Washington, DC 20234

Thomas J. McGee University of Maryland Molecular Physics Bldg.

College Park, MD 20742

Joe V. Michael NASA/Goddard Space Flight Center

and Catholic University of America

Code 691 Greenbelt, MD 20771

David F. Miller

Battell e - Columbus Laboratory

505 King Avenue

Columbus, OH 43201

Mario J. Molina University of California

Department of Chemistry

Irvine, CA 92717

Hiromi Niki Scientific Research Lab. Ford Motor Company

P. 0. Box 2053 Dearborn, MI 48121

Robert J. O'Brien Portland State University

Department of Chemistry

P. 0. Box 751 Portland, OR 92707

H. Edward O'Neal

San Diego State University

Department of Chemistry

San Diego, CA 92115

David A. Parkes

Shell Research B.V. Badhuiswes 3 Amsterdam N Postbus 3003 The Netherlands

Robert Allen Perry

ERL/N0AA

Department of Commerce

Radio Building, Room 3522 Boulder, CO 80302

A. R. Ravishankara

Georgia Institute of Technology

Applied Science Lab., EES Georgia Tech. Atlanta, GA 30332

Keith Schofield

ChemData Research

P. 0. Box 40481 Santa Barbara, CA 93103

John H. Seinfeld California Institute of Technology

Pasadena, CA 91125

Donald H. Stedman University of Michigan

c/o NCAR, P. 0. Box 3000 Boulder, CO 80303

104
Louis J. Stief NASA/Goddard Space Flight Center

Code 691 Greenbelt, MD 20771

Fred Stuhl NOAA/ERL, Aeronomy Laboratory

325 Broadway, R448 Boulder, CO 80303

Wing Tsang

National Bureau of Standards

A145, 222 Washington, DC 20234

Peter Warneck Max-Planck-Institut fur Chemie (Otto-Hahn-Institut)

65 Mainz, Saarstr. 23 Germany

Robert Tony Watson

Jet Propulsion Laboratory

California Institute of Technology

Building 183-601 4800 Oak Grove Drive Pasadena, CA 91103

Karl Westberg

The Aerospace Corporation

P. 0. Box 92957

Los Angeles, CA 90009

Michael Whitbeck

General Motors Research

Environmental Science Department

Warren, MI 48090

Gary Z. Whitten

Systems Applications, Inc. 950 Northgate Drive

San Rafael , CA 94903

105

Subject Index Author Index

Aldehydes, 27 Alkenes (see ozone-al kene)

Alkoxy radicals, 20,34,51,62,65

Alkyl nitrate, 77 Aromatic compounds, 85,92,93,95

Creigee mechanism, 7,21,22

Creigee zwitterion, 64

Dialkenes, 19

Formaldehyde, 46,47

Formic Acid, 47,61

Formyl radical, 47

H0 2 , 34,99

HS0 4 radicals, 34 Hydroxyl radical, 7,15,31,47,60

74,77,85,92,93,95,99,100

Monoterpenes , 19

Nitrogen oxides, 71,74,78

N0 3 radicals, 34

0 atom, 32,89

OH-olefin, 9,20 Olefins (see ozone-al kene)

0' Neal -Bl umstein mechanism, 22 Oxygenated hydrocarbons, 27 Ozone, 7,20,23,78,89

Ozone-al kene, 7,12,15,20,22,23,61

Peroxyacetyl nitrate (PAN), 35,78

Peroxyacyl nitrate, 35,77

Peroxyalkyl nitrates, 77 Peroxy radicals, 64,65

Sulfur oxides, 99

S0 2 , 99,100

Atkinson, Roger, 15,92

Batt, L. , 21 ,62 Benson, Sidney W. , 64

Calvert, Jack G. , 99 Carter, William P. L, , 20,65,77,78

Cox, Richard A. , 65,71 ,100

Doty, Richard M. , 73,74

Dodge, Marcia C. , 21 ,78

Golden, David M. , 51 Green, Patrick J. , 74,93

Heicklen, Julian, 20,23,46,47

Hendry, Dale G. , 77,85

Jeffries, Harvey E. , 100

Lloyd, Alan C. , 27

Miller, David F. , 100

Niki, Hiromi, 7

O'Brien, Robert J. , 74,93

O'Neal , H. Edward, 22

Ravishandara, A. R., 47,95,101

Tsang, W., 22,64,79

Whitten, Gary Z. , 22

107
NBS-114A (REV. 0-78)

U.S. DEPT. OF COMM. BIBLIOGRAPHIC DATA

SHEET

1. PUBLICATION OR REPORT NO. SP 557

2. Gov't Accession No, 3. No.

4. TITLE AND SUBTITLE Chemical Kinetic Data Needs for Modeling the Lower

Troposphere

5. Publication Date August 1979

$, Perforu*iGg Organization Code

7. AUTHOR(S)

J. T. Herron, R. E. Huie, and J. A. Hodgeson, editors

8. Performing Organ. Report No.

9. PERFORMING ORGANIZATION NAME AND ADDRESS

NATIONAL BUREAU' OF STANDARDS

DEPARTMENT OF COMMERCE

WASHINGTON, DC 20234

10. Project/Task/Work Unit No.

11. Contract/Grant No.

12. SPONSORING ORGANIZATION NAME AND COMPLETE ADDRESS (Street, City, state, ZIP) Sponsored in part by

Environmental Protection Agency

Research Triangle Park, NC 27711

13. Type of Report & Period Covert Final

14. Sponsoring Agency Code

15. SUPPLEMENTARY NOTES Library of Congress Catalog Number: 79-600125

|

| Document describes a computer program; SF-185, FIPS Software Summary, is attached.

16. ABSTRACT (A 200-word or less factual summary of most significant infomtation. If document includes a significant bibliography or

literature survey, mention it here,) the proceedings of a workshop on chemical kinetic data I ov/er troposphere, held at Reston, Virginia, May 15-17, 19]

y the Environmental Protection Agency and the National i ussed on six key problem areas in tropospheric

olefins with hydroxyl radicals and ozone, reactions of I reactions, reactions of oxides of nitrogen, reactions of

j reactions of oxides of sulfur.

This is a report of

needs for modeling the T

The meeting, sponsored b\

Bureau of Standards, foci

chemistry: reactions of

aldehydes, free radical

aromatic compounds, and

The report includes

work, review papers, dis

and an attendance list.

a summary and list of major recommendations for further

cussion summaries, contributed comments, recommendations,

J

17. KEY WORDS (six to twelve entries; alphabetical order; capitalize only the first letter of the first key word unless a proper name;

separated by semicolons) Aldehydes; aromatics; chemical kinetics; data needs; free radicals; modeling;

NO ; olefins; SO ; troposphere.

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