Introducing Atmospheric Reactions A Systematic Approach for Students N. Colin Baird University of Western Ontario, London, ON, Canada N6A 587 As cmvirnnnicntal srienre incrensei in importanre hoth to students and to ~ r o f ~ ~ s s i o n athe l s , topic of ntmi,sphenc chemistry also grows in relevance-even finding its way into introductory textbooks for general ( 1 )and for environmental chemistry (2).Unfortunately, air is a complex reaction medium: over 200 reactions occur simultaneously in polluted urban air (3).Consequently, there is a tendency among educators either simply to ignore the mechanisms of atmospheric reactions and simply quote the overall result or to present a n arbitrarily chosen small numbw of reactions a s representative. In what follows, an attempt is made to outline the dominant reactions that occur i n air, to show that the relative importance of such processes can be rationalized using elementarv urinciules.. and to illustrate a few "rules" usine which tjle'predominant fate of many pollutants ran be prrdieted. In essence. the half-duzen or so reaction t .v.~ c of s greatest importance to the chemistry of the lower atmosphere have been deduced from the verv large set of reackons in air discussed by Finlayson-pitti and-~itts(3a) and are presented here. ~~~~~~~~
.
-
Organizing Principles for Air Chemistry As has been noted many times, the Earth's atmosphere is a n oxidizing environment owing to the great concentration of Oz present. Almost all the gases released into airwhether "natural" substances or pollutants-are eventually oxidized; e.g. CH, + 20,
+ CO, + 2HzO
Intact 0, molecules do not directly react with the original gases! The homogeneous oxidation of covalent molecules in air inevitablv occurs bv a seauence of free radical reactions. The "activation" of the original molecules occurs either when thev absorb sunlight - and undergo .. photo. chemical decomp(~sitim,or in the more t'rrqucnt case when thev are attackvd hs the hvdroxyl t'rer radical,' OH' lor in rare cases by ozone). Examules of ~hotochemicaldecom~ositionmainly involve t i e carbdnyl group; most impokant are the Hldehvdes H?CO and CHaCHO which are released in automobile exhaust and which are produced as intermediates in the atmospheric oxidation of other pollutants, for example: sunlight W
H&O
H'
+ HCO'
-
(A fraction of aldehvdes eive molecular products such a s " Hz.) Such reactions increase the concentration of free radicals in air. which i n turn a e a t l v increases the rate of the overall oxidation reaction; and is the driving force in photochemical smog episodes in Los Angeles, Mexico City, and many other warm, sunny cities. The feasibility for sunlight to produce free radicals from other gases can be easily de'we adopt here the standard convention that a superscript dot place after a formula indicates that the species is a free radical; whereas, one place directly above an elemental symbol indicates that the unpaired electron is localized on that atom.
termined by cnlcul:~tingthe A l I of the dissociation reaction and noting that the sunl~ghtcutoft' wavelrnfirh of about 290 nm corresponds to an enwhy of413 k.1 mol ' = 99 kcal mol-'. Much more common than photolysis of closed-shell molecular pollutants is their reaction with OH', a very reactive species whose concentrati~nin air is miniscule (about 5 x lo6 molecules per mLj but without which most pollutants would not be oxidized a t a measurable rate. Reactions involving the hydroxyl radical readily occur provided that their activation energies do not exceed a f e a kilocalories per mole; thus, endothermic reactions involving OH' are not important because their rates are so slow. The lowest activation energy-and hence fastest -reaction for a molecule involves addition of OH' to a n atom which is participating i n a multiple bond, e.g. a n alkene: HzC=CH2 + OH' + H,CCH,OH The site of the unpaired electron is, thereby, altered from the oxygen atom in OH' to the atom of the double bond to which the oxveen . does not attach. In order for the reaction to be exothermic, the newly-formed single bond to oxygen must be stronger than the oi bond which is destroyed: this condition is n i t fulfilled if h e atom attached to ;xy;en is also an oxygen atom (since 0-0 bonds are weak), or if the pi bond i s that of a carbonyl group (since i t is stronger than is a CO single sigma bond). Hydroxyl radicals do not add to any fully-oxidized product, such a s COz, SO3, HNO3, or N205, since the multiple bonds therein are so strong. Similarly OH' addition does not occur in nitrous oxide, NzO, because of the high bond strengths in the reactant and the weak N-0 bond that would result. If addition of OH' is not feasible due to lack of a pi bond or for energetic reasons, then i n hydrogen-containing molecules X-H the process of H atom abstraction by OH' usually occurs because its activation energy is usually only slightly greater than that typical of addition: X-H + OH'
+ X'+ H,O
The strength of the H-OH hond thereby formed exceeds that of almost all X-H bonds (a notable exception being the H-F linkage) and so no large harrier to the reaction is imposed by its energetics. According to the reasoning above, t h e first products formed by the activation of a molecular gas in air are free radicals. In most cases, the next reaction incorporates diatomic oxygen into the system a s a whole. One common mechanism by which this occurs is the abstraction by 0 2 of a hydrogen atom from the free radical to form the hydroperoxv radical, HOW. Since the H - 0 0 bond is not strong, ?),can nl,itract hydrogen from radicals only if this procris i~llowjthe simultaneous formation of a nea. pi bond within the original radical, thus compensating energetically for the loss of the bond to hydrogen. Usually the new bond forms bv combinine the electron that had been bonded to H with the unpairLd electron i n the free radical; thus, the hydrogen to be abstracted must be originally bonded to a n Volume 72 Number 2 February 1995
153
transfer of a n oxygen atom to NO'; thus the peroxy radical oxidizes nitric oxide to nitrogen dioxide: Does radical have a peroxy bond?
1
,
NO;
~
-C
+ NO' + OH' + NO; XOO' + NO' + XO' + NO;
HOO'
radical
Is it of type No F
+ O-centered
aldehyde +radical
H ?~
C-csplits (i Rfstabilizing)
Can H removal from H-X convert X-0' (or X=O) to X=O (or XEO)?
Olgb&d
HOCI '+ non-radical
By this process, most of the NO' in urban atmospheres is oxidized to NO;, and most hydroperoxy radicals are converted to hydroxyl radicals. Many oxygen-centered free radicals XO' of interest conto form tain a hydrogen atom that can be abstracted by 0% HOO' and a molecule which contains a double bond to the oxygen. In other XO' radicals, there is a chain of carbon atoms attached to the oxygen; scission ofthe carbon-carbon bond nearest the oxygen produces an aldehyde and a carbon-centered free radical. Because this process is only mildly endothermic (about +12 kcal mole-') when the carbon chain is unbranched, substitution a t the radical site in the product by an oxygen group or several alkyl groups lowers the activation to almost zero; in such circumstances, the bond scission process occurs spontaneously and is faster than hydrogen abstraction by 0 % .Specific examples of these XO' reactions are discussed in the next sections. The organizing principles developed above are summarized as "decision trees" in Figures 1and 2.
Applications The reaction principles developed above can be applied readily to deduce the fate of many atmospheric pollutants. Consider for example the reactions of sulfur dioxFigure 1. The fate of airborne free radicals, under conditions with significant ide gas emitted into the atmosphere if it is oxidized by nitric oxide present and beforeradical + radical reactions become important. homogeneous processes.2 Since so2does not absorb sunlight and contains no hydrogen, it is stable until activated by addition of O R to the sulfur atom: atom bonded to the radical site. For example, if a gaseous methanol molecule is activated by abstraction of a methyl group hydrogen to yield the radical CHzOH', yes Gas eventually returned results in the then abstraction of the hydroxyl H by 0% ) to Earth's Burface conversion of the CO single bond to a double bond: soluble or fully 0, ad& to radical site to produce perom radical
H2CO-H + 0%+ H2C=0 + HOO' Similarly, the HCO' free radical produced by photolysis of formaldehyde is converted to carbon monoxide by abstraction of its hydrogen: H-&O
+
0%+
ce=o@+ HOO'
If hydrogen abstraction is not feasible because a pi bond cannot be formed, then 0% generally adds to the free radical a t the site of the unpaired electron, creating a peroxy free radical X 0 0 ' analogous to the hydroperoxy of the radical HOW. An example is the reaction with 0% methyl free radical (created by the abstraction of a H by OH' from methane): ,
CHj
+
0%+ CH,OO'
Atrivial example is the reaction of Ozwith atomichydrogen. Compared to the free radicals produced initially by photolysis or OH' attack, the peroxy free radicals are relatively long-lived. They generally cannot abstract hydrogen because the OO-H bond is a weak one and thus such reactions are endothermic and would have large activation energy barriers. I n urban atmospheres and other air masses containing the stable free radical nitric oxide, NO', the usual reaction of peroxy radicals is the
D o e s gas photodecompose
'PI1: in sunlight?
Do gea moledes
that OH' have multiple can add bonds
>
OH'addition occurs
2~uch of the SO, is actually oxidized by heterogeneous procGm i. inert in esses involving water droplets; the mechanism, which involves troposphere; will rise ozone and hydrogen peroxide as the oxidizing agents, is beyond to stratosphere the scope of the present discussion of homogeneous reactions. Figure 2,The fate of gases emitted intoair, 154
Journal of Chemical Education
Freer!Asare prcdueed
Finally, the four HOO' free radicals produced in the above reactions eventually oxidize the same number of nitric oxide molecules:
+
4(HOO'
This free radical then reacts with diatomic oxygen by abstraction of the hydrogen, since this allows conversion of S-0 to S=O:
NO' + OH'
+ NO;)
If all 10 reaction steps above in the methane oxidation sequence are added together, the overall reaction is seen to be
sunlight
CH,
The sulfur trioxide so formed is fully oxidized; it quickly combines with water vapor to form droplets of sulfuric acid aerosol, HzSO,(aq). A second example is the oxidation of methane, CHI, which is emitted in vast quantities into the atmosphere as a result of both "natural" and anthropogenic activities. Since methane does not absorb sunlight and contains no pi bonds, OH' abstracts a hydrogen atom from it:
+
CH,
+
OH'
CH; + H20
(Since the activation energy here is several kilojoules, this reaction is relatively slow; on average a methane molecules exists in air for a decade before it is activated. Other alkanes react more quickly since the activation energies for abstraction are near zero.) As predicted from our principles, the methyl radical reacts with diatomic oxygen; since no pi bond can form, the oxygen adds to the carbon atom since it is the radical site:
The methyl peroxy rndlcal then oxidizes nitricoxide, pnlducing the methoxy free radical and nitrogen dioxide:
+
CH300'
NO'
+ CH30'
+ NO;
In the electronic structure of the methoxy radical, the unpaired electron is sited a t the oxygen atom, and there are hydrogen atoms bonded to the carbon atom; thus, the next reaction is abstraction by 0, of a hydrogen and the simultaneous creation of a carbon-oxygen pi bond, i.e. formaldehyde:
+
CH30'
0,
'HzC=O + HOO'
Formaldehyde is a natural constituent of air because it is an intermediate in the oxidation of methane. As previously discussed, in air it photodissociates in sunlight to produce free radicals HzCO
+ sunlight + H' + HCO'
As discussed above, both these radicals react with diatomic oxygen to produce the hydroperoxy free radical. H'
+
H-&O
+ Oz -t Ce 'o@ + HOO'
OZ -t HOO'
+ 5 NO' + SO2
In a sense, nitric oxide is co-oxidized with methane-and with many other substances-in atmospheric reactions. This mechanism corresponds to the ambient temoerature "burning" of hydrocarh&s in air. Further Reactions of Free Radicals The "rules" quoted above are appropriate to the background concentrations of free radicals found in most air masses-in particular to those which are neither particularly "clean" nor particularly "dirty". In episodes of high ~ollutionin manv urban areas. the concentration of the tree radicals becomes so great that other processes, especially radical-radical reactions, achieve importance. Wjlen two species with unpaired electrons combine, generally a closed-shell molecule is produced. For example, when late in the day of an episode of urban smog the concentrations of both the hydroxyl radical and nitrogen dioxide become relatively high, they combine with each other to produce one of the final products of the process, nitric acid: OH'
+
Ce-O@
+
OH'
+ H-O-C=O
Clearly, diatomic oxygen will abstract the hydrogen from this free radical since a n additional CO pi bond is thereby formed; in this way, the fully-oxidized gas carbon dioxide is finally obtained: HOCO' + OZ -, O=C=O + HOO'
NO; -, HNO,
If the combination of the two radicals yields a species with a chain of three or more oxygen atoms, then one or two O2 molecules split out since 00 single bonds are relatively weak and unstable; for example two hydroperoxy radicals combine to yield hydrogen peroxide and diatomic oxygen: 2 H00'
+ [HOOOOH] + H,Oz +
OZ
Nitrogen dioxide, NO; , although a free radical is relatively stable in air. As illustrated above, it can participate in radical-radical reactions. Alternatively, when exposed to sunlight it can photodissociate by detaching a n oxygen atom: sunlight
NO;
A
NO'
+
0
Since the O-NO bond energy is 306 kJ mol-', the maximum wavelength for this process is about 391 nm, i.e. in the UV-A region. The oxygen atoms produced here quickly combine with diatomic oxygen to produce ozone: 0
The carbon monoxide obtained here is also a natural intermediate of methane oxidation; since the second pi bond in it is weak and since CO does not absorb sunlight, OH' adds to the carbon atom:
+ HzO + 5NO; + 2 OH'
CO,
+
0,
" o3
This sequence is the principal route by which ozone is produced in both clean and polluted air. Photochemical Smog: The Problem of "Urban Ozone" In relatively "clean" air, the reactions described above do not produce pollutants in sufficient concentration to be troublesome to the health of living matter. In contemporary civilization, though, huge volumes of very reactive hydrocarbons and aldehydes, and huge volumes of nitric oxide, NO', are emitted each day into the air of urban areas. As a consequence, the oxidation reactions proceed much more rapidly and produce much greater quantities of noxious products than in "clean" air. This is particularly true in cities that are especially sunny and warm, and whose Volume 72 Number 2
February 1995
155
STEP RHC = CHR
into u r b a n a i r from vehicle exhaust, particularly if the fuel contains a n alcohol. I n a sunlit atmosphere, t h e aldehydes decom ose photochemically (reaction 6) t o produce two new free radicals R' and HCO'. This step i s a critical one i n t h e photochemical smog process, since i t increases t h e free radical concentration i n air and makes the atmosphere much more reactive than i s "clean" air. Although the same sort of process occurs with methane, i t s effect there i s much less since methane m o l e c u l e s on a v e r a g e r e q u i r e years before undergoing successful OH' attack, whereas alkenes react with i t i n a matter of minutes. Inclusion of t h e conversion of HCO' to carbon monoxide (reaction 7b) and the addition of 0 2 to R' (reaction 7a) and the subsequent oxidation of NO' by the peroxy radical (reaction 8 a ) a n d by t h e HOW radicals produced i n reactions 5 and 7b yields the following overall reaction: RHC=CHR + 6NO' + 6 0 2 +sunlight
B
RHC = 0
+
RH~oH
1
Oz RHC = 0
+
A+
4% +
HOW
NO'
N 4
RO'
HOO'
HCO
7 a.b
ROO'
+
OH'
4
CO
NO'
+ NO;
Figure 3. The reaction sequence for a prototypical alkene, RHC=CHR,where R is a short alwl chain,
air masses can be immobilized temporarily by geographic constraints such a s mountains. The greatest mass of reactive hydrocarbons are generally those emitted by motor vehicles-either by evaporation of their fuel or by emissions of incompletely-burned fuel through their tailpipes. The most reactive class of hydmcarbons are the alkenes, followed by aromatics and alkanes. As a n illustration of the photochemical smog which can be generated in urban air, the reaction sequence for a pmtotypical alkene RHC=CHR, where R is a short alkyl chain, is illustrated in Figure 3. As anticipated from the reaction principles discussed above, the reaction sequence is: 1. Addition of OH' to a carbon of the douhle bond 2. Addition of O2 to theother carbon of the original double
T h e alkoxy r a d i c a l s RO' s u b sequently react with O2 t o produce additional aldehyde molecules by loss of hydrogen attached to t h e alpha carbon; their photochemical decomposition yields a carbon chain with one less member t h a n was present i n R. The degradationofthe alkylchaincontinuesuntilitis all converted to carbon monoxide, with a corresponding increase i n the number of free radicals i n the air each time t h a t a carboniseliminated. As noted previously, the increase i n the free radical concentration does not continue unabated in air, because other reactions become important when radicals abound. For example, when [OH'] becomes high, some of the aldehyde is degraded by OH' abstraction of the carbonyl hydrogen rather than by photolysis:
bond (since it is the site of the unpaired electron) 3. Oxidation afNO' by 0 atom transfer from the peroxy radi-
cal Instead of hydrogen abstraction by 0 2 from the alkoxy radical A, spontaneous decomposition occurs by scission of the C-C bond to produce the aldehyde RHC=O and the radical RHCOH'; this process is almost thermoneutral and thus i s of low activation energy since a carbonyl CO pi bond is not much weaker than a CC single bond and because the product radical is stabilized by the -OH group. In the subsequent reaction of the new radical, O2 abstracts the hydroxyl hydrogen to produce a second RHC=O aldehyde molecule (reaction 5). Thus, the original alkene RHC=CHR produces two aldehyde molecules RHC=O, each half the size of the original hydrocarbon. Aldehydes also a r e emitted directly 3These products occur more abundantly than do the alternatives He and RCO'; for simplicity, the alternative products are ignored
herein.
156
Journal of Chemical Education
The R-C=O radicals, as anticipated, add Oz to produce peroxy free radicals which then oxidize NO' if the latter is abundant:
The RCOp radicals decompose spontaneously by R-C bond scission to yield COz; this process is exothermic and therefore fast since the pi bond in O=C=O is much stronger than is a C-C single bond.
Once most of the nitric oxide has been oxidized to NO; and the radicals concentrations have vastly increased, that is, the situation typically found during afternoons in polluted urban atmospheres-radicals such as R-C=O react instead by addition to other radicals such as NO;
. .
sential orocesses in t r o ~ o s ~ h e rchemistw: ic .. the manv reactions ignored in t h i s sch(:mi, ;in! thwc that are rever.;ible ~ l "a vlittle and. rhu>., "~ i r l dno net uroducts. or [hose whirh . part in the overall reactions and lead only to products of minor importance. These principles have been used for several years by the author in an introductory environmental chemistry course for science majors, all of whom have had a t least a rudimentarv introduction to oreanic chemistw and es~eciallv " to Lewis structures. Students seem capable of successfully usine the ~ r i n c i ~ lto e sdeduce reaction mechanisms for the oxidation of gaseous trace substances such a s HzS, ethene, simple alkanes, etc. "
The prototypical radical product of this type corresponds to R=CHz, and is known a s PAN = peroxyacetylnitrate; it is a powerful eye irritant:
A
Literature Cited
PAN
Conclusions The organizing principles for homogeneous reactions in atmospheric chemistry are successful in isolating the es-
3. 1s) Finlayson-PiftP, B. J.;Pitts,J. N.,Jr.Afmofpheheii Chemistry: WiIey-Intefefefeience, John Wiley and Sons: New York, 1986. ibl Heicklen, J.Atmospheric C h ~ m k l r y ; Academic Press: New York. 1976. (el Seinfeld, J. H. Atmospheric Chemistry and Phyricr ofAir Pnllulion; John Wiley and Sons: New York, 1986.
Volume 72 Number 2
February 1995
157
