Atmospheric Fate of CarbonylOxidation Products Originating fromr-Pinene and ∆3-Carene:Determination of Rate of Reactionwith OH and NO3 Radicals, UVAbsorption Cross Sections, andVapor PressuresM A T T I A S H A L L Q U I S T , * , †

I N G V A R W A N G B E R G , ‡ A N DE V E R T L J U N G S T R O M †

Department of Inorganic Chemistry, University of Goteborg,S-412 96 Goteborg, Sweden, and Physikalische Chemie/Fachbereich 9, Bergische Universitat-GesamthochschuleWuppertal, Gaussstrasse 20, D-42097 Wuppertal, Germany

Large yields of dicarbonyl compounds have been foundin the atmospheric oxidation of R-pinene and ∆3-carene.These terpenes are emitted in large quantities by biogenicsources, and it is important to know the fate of their reactionproducts. In this investigation, ultraviolet and infraredabsorption cross sections, vapor pressures, and rate coef-ficients for hydroxyl and nitrate radical reactions havebeen determined for the main product in each case, i.e.,3-acetyl-2,2-dimethylcyclobutaneacetaldehyde (pinon-aldehyde) and 2,2-dimethyl-3-(2-oxopropyl)cyclopropaneacet-aldehyde (caronaldehyde). Photolysis lifetimes at noonon July 1 at 50° N using a photolysis quantum yield of 1 are3.3 h for pinonaldehyde and 5.8 h for caronaldehyde. Theinfrared absorption cross sections obtained (base e) were(8.08 ( 0.32) × 10-19 and (1.04 ( 0.05) × 10-18 cm2

molecule-1 for pinonaldehyde and caronaldehyde at 1725.4and 1741.5 cm-1, respectively. Vapour pressures deter-mined by Knudsen effusion measurements are 5.1 and 3.0 Paat 298 K. The rate coefficients obtained using the relativerate technique for reaction with OH radicals were (8.72 (1.14) × 10-11 and (1.21 ( 0.36) × 10-10 molecules cm-3

s-1 for pinonaldehyde and caronaldehyde and for the cor-responding NO3 reactions were (2.35 ( 0.37) × 10-14 and(2.71 ( 0.15)× 10-14 molecules cm-3 s-1. The vapor pressuresare too high for homogenous nucleation or directcondensation to take place in the atmosphere. The dominantgas-phase removal processes will be reaction with OHradicals and possibly also photodissociation.

IntroductionEmission of monoterpenes is a major source of organiccompounds in the troposphere. Monoterpenes constituteabout 11% of the total global amount of naturally emitted,volatile non-methane organic compounds (1). Locally, theemission of monoterpenes may even dominate the volatile

organic load into the atmosphere. Kinetic investigations haveshown that major sinks for monoterpenes in the atmosphereare reactions with hydroxyl radicals, nitrate radicals, andozone (2). Some of the products from the oxidation ofmonoterpenes are believed to contribute to aerosols of naturalorigin found in the atmosphere (3). Aerosol formation in theatmosphere is an important issue since particles may act ascloud condensation nuclei, may affect visibility, and couldlocally be troublesome from a health aspect. Aerosol particlesmay also catalyze heterogeneous reactions and removepotential ozone-forming compounds from the gas phase (4).

Product information, e.g., from smog chamber experi-ments is available for several monoterpene-oxidant systems(3, 5-10). The mass balances reported are in most cases farfrom closed. The mass deficiency is frequently explained bythe formation of aerosols in the reaction mixture, removingmaterial from the gas phase. Part of this aerosol formationcould be caused by the high concentration of reactants thatis often used in such reaction systems. The major gas-phaseproducts found in monoterpene-oxidant systems are varioustypes of carbonyl species, originating from the oxidative attackon a double bond. Two of the dominant monoterpenesemitted from coniferous forest are R-pinene and ∆3-carene(11). Studies of the oxidation of these two terpenes underatmospheric conditions in static reactors show large yieldsof 3-acetyl-2,2-dimethyl-cyclobutaneacetaldehyde (knownas pinonaldehyde) and 2,2-dimethyl-3-(2-oxopropyl)cyclo-propaneacetaldehyde (by analogy, here known as caronal-dehyde) (3, 5-10). Both products and starting materials areshown in Figure 1. These 10-carbon atom dicarbonylcompounds are expected to have a much lower vapor pressurethan the corresponding monoterpene. They may thereforebe suspected of contributing to natural particle formation inthe atmospheres and pinonaldehyde has been identified inambient aerosols (3).

Several gas-phase oxidation routes are available if pinon-aldehyde or caronaldehyde are to be removed from thetroposphere. The ketone and aldehyde groups make themprone to photodissociation. Other possible, competitivechannels are reaction with hydroxyl radicals and, duringnighttime, reaction with nitrate radicals. The aim of this workwas to determine the relative importance of different removalprocesses of pinonaldehyde and caronaldehyde. The workincludes determination of OH and NO3 radical reaction ratecoefficients, ultraviolet (UV) absorption cross sections im-portant for photodissociation, and vapor pressure data neededfor assessment of the aerosol formation capability of the twodicarbonyls.

Experimental SectionVapor Pressure Measurements. The vapor pressure mea-surements were made by using a Knudsen effusion apparatus.The substance, the vapor pressure of which is to bedetermined, is placed in an aluminium container equippedwith a copper foil lid, sealed by an o-ring, and held down bya nut. The lid is supplied with a circular effusion orifice. Twonominal orifice diameters were used in this investigation, 0.3and 0.9 mm, in foils of 0.05 and 0.1 mm thickness, respectively.The complete cell is placed in a copper block inside a vacuumchamber that is pumped by a 120 L s-1 diffusion pump. Thepressure in the vacuum chamber is estimated by using aPenning gauge. The copper block is cooled by a two-stagePeltier element unit, set at a constant cooling power suitablefor the measurement to be made. The temperature of theblock is held constant to within (0.05 K by controlled heating.Pt-100 thermometers are used, both for the temperaturecontroller and for the actual temperature measurement of

* To whom correspondence should be addressed. E-mail:hallq@inoc.chalmers.se; fax: + 46 31 7722853.

† University of Goteborg.‡ Bergische Universitat-Gesamthochschule Wuppertal.

Environ. Sci. Technol. 1997, 31, 3166-3172

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the container. Measurements for the dicarbonyls were madein the temperature range from 255 to 276 K, where the vaporpressures fell between 0.01 and 0.5 Pa. (1 Pa ) 0.01 mbar).This means that the pressure in the vacuum chamber wasalways at least a factor of 100 lower than in the effusion celland that the mean free path in the cell was always considerablylonger than the orifice diameter but not necessarily greaterthan the cell dimensions. Experiment times between 2 hand 2 day were used, resulting in weight losses between 0.2and 5 mg. Experiment times were long as compared with thetimes for the transients at the beginning and end of anexperiment. The vapor pressure may be calculated from theweight loss and the nominal orifice dimension or via acalibration procedure for the orifice.

IR Absorption Calibration. Accurate absorbance-concentration calibrations of pinonaldehyde and caronal-dehyde in the infrared (IR) were necessary for the concen-tration measurements in the determination of UV absorptioncross sections. Measurements of IR absorption cross sectionswere made using the 0.480 m3 borosilica glass reactordescribed below, in the UV absorption measurements section.Due to the comparatively low vapor pressures of the dicar-bonyls, these compounds are difficult to transfer in aquantitative manner to the gas phase, and they are lost byadsorption to the reactor walls. A special procedure wastherefore developed to handle these problems. The mea-surements of the absorption cross sections were performedat low pressure, <10 Pa, and at 300 ( 1 K. An experimentconsisted of first evacuating the reactor and then introducinga sample of the dicarbonyl. Then, five FTIR spectra, eachconsisting of 10 co-added scans, were collected, and thereactor pressure was noted in the middle of each collection.The time needed for each spectrum was 17 s, and anexperiment lasted approximately 5 min. The absorbance ofthe dicarbonyls was always lower in subsequent spectra dueto loss by adsorption to the reactor walls. Pressures weremeasured by using a Baratron type 128 pressure gauge, witha resolution of 0.01 Pa. Reactor leak rates were also measuredbetween each experiment. The rates were always very low,3-8 × 10-3 Pa min-1, but significant for the measuredpressures and thus had to be accounted for. In experimentswith caronaldehyde, the presence of methanol and dichlo-romethane, solvents originating from the synthesis, alsocontributed to the pressure. Using known absorbance crosssections for these compounds, concentrations and therebytheir contribution to the total pressure were calculated andcorrected for. Each experiment yielded a set of absorbancesand dicarbonyl concentrations from which the absorptioncross section could be calculated.

To test the accuracy of the pressure measurementsdescribed above, known amounts of liquid R-pinene wereintroduced into the evacuated chamber, and the pressuredifference was measured. The R-pinene samples were rapidly

and completely evaporated into the reactor, and accordingto the FTIR measurements, no wall loss was occurring. Thus,the increase in pressure was due only to the vapors of R-pineneand to leaking. Excellent agreement between the measuredpressure increase, after compensation for leaks, and theexpected pressure increase as calculated from the introducedvolumes of liquid R-pinene were obtained.

UV Absorption Measurements. The experimental setupused for UV spectroscopic measurements is located inWuppertal and has been described in detail elsewhere (12).However, a short presentation will be given here. The 0.480m3 borosilica glass reactor has a length of 3 m. The reactoris evacuable, and most of the spectra were obtained at lowpressure. Two long-path mirror systems, one for IR and onefor UV absorption spectroscopy with optical path lengths of51.6 m and 39.1 m, respectively, are mounted inside thereactor. The UV system consists of a SPEX 22-cm mono-chromator combined with a PAR 1412 diode array detector,yielding a spectral resolution of 0.6 nm. IR spectra weremeasured with a Nicolet 520 FTIR spectrometer working at1 cm-1 resolution. The IR system was used to measure theabsolute concentration of the investigated compound for eachUV spectrum. A UV spectrum was obtained by first measuringthe intensity I0 at each diode without any substance presentand then introducing the dicarbonyl compound and recordinga new set of intensities, I. The absorbance, ln (I0/I), for eachdiode was then used to calculate the absorption cross sectionaccording to Beer-Lambert’s law. Each UV spectrumconsisted of 6 or 7 co-added intensity measurements, andthe time required for collection of each spectrum was between108 and 112 s. Several UV spectra were recorded at differentconcentrations and normalized, and the mean absorptioncross section for each diode in the array was then calculated.A calibration to set the wavelength scale on an absolute basiswas done, using the two mercury lines at 296.73 and 312.57nm.

Relative Rate Experiments. The relative rate experimentswere made using a reactor at the University of Goteborg. Thisexperimental setup and the procedures used in the relativerate experiments have been described in refs 13 and 14 andonly brief information will be given here. The borosilica glassreactor has a volume of 0.153 m3 and a length of 2 m. It isprovided with multi-pass optics connected to a MattsonPolaris FTIR spectrometer that is equipped with an MCTdetector. Path lengths of 24, 32, or 40 m were used in theexperiments. The spectral resolution used was 1 cm-1. Thelight source for photolysis in the OH reaction experimentswas a bank of three, 2 m long, 100 W Philips UVA fluorescentlamps, placed approximately 10 cm above the reactor.

The relative rate coefficients for reaction of the dicarbonylswith NO3 or OH were determined in experiments usingpropene as the reference compound. In such an experiment,the radical reacts with the dicarbonyl, with the unknown ratecoefficient k1, and also with the reference organic compoundwith the known rate coefficient kpropene. For the dicarbonylcompounds there is also a first-order loss channel due toadsorption to the wall, process 3. For the OH experimentsthere may also be a first-order loss through photodissociation.In the present case, this process was negligible, compared tothe wall loss and was therefore excluded from the analysis:

Provided that reactions 1-3 are the only loss channels for theorganic compounds and that neither the dicarbonyl nor the

FIGURE 1. Structural relation between r-pinene and pinonaldehydeand between ∆3-carene and caronaldehyde.

dicarbonyl + NO3 or OH 98k1

products (1)

propene + NO3 or OH98kpropene

products (2)

dicarbonyl 98kw

dicarbonyl(wall) (3)

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propene is re-formed in any process, then the decrease inconcentration for the two organic compounds is describedby eq I.

Subscripts 0 and t correspond to concentrations at thebeginning of an experiment and at time t, respectively.Absolute concentrations may be substituted by IR absor-bances if Beer-Lambert’s law holds. The linearity betweenIR absorption and concentration was confirmed by a calibra-tion procedure for each compound used in the experiments.The rate coefficient, k1, was obtained from the slope of a plotof ln([dicarbonyl]0/[dicarbonyl]t) - tkw vs ln([propene]0/[propene]t) combined with a literature value of kpropene of 9.3× 10-15 and 2.63× 10-11 cm3 molecule-1 s-1 for NO3 (15) andOH (2), respectively. Wall loss of the carbonyls, kw, wastypically (4-7) × 10-5 s-1 and was determined in separateexperiments. The reference, propene, is not lost to the reactorwalls.

Hydroxyl radicals were generated by photolysis of CH3-ONO in presence of NO according to the following reactions:

The experiments were performed in synthetic air. Nitrateradicals were generated by the thermal decomposition of N2O5

at room temperature. These experiments were conductedwith nitrogen as the matrix gas.

Chemicals. Synthetic air (20% O2), nitrogen (g99.995%),propene (>99.8%), and NO (2.0% in nitrogen) were purchasedfrom AGA Gas AB. The naphthalene used was Merckscintillation grade. N2O5 was prepared by the reactionbetween O3 and NO2, with O3 in excess, in a flow system. Theproduct was trapped and stored at 195 K. CH3ONO wassynthesised by treating an aqueous solution of NaONO andmethanol with H2SO4 (16). Pinonaldehyde and caronaldehydewere made by ozonolysis of the corresponding terpene (17).Both pinonaldehyde and caronaldehyde were analysed usingGC-MS and NMR technique. The sample of pinonaldehydewas analyzed by two-dimensional (1H/13C) NMR and DEPTspectra and found to be of 95% purity. A similar purity wasobserved by GC-MS. The MS spectrum agreed with theliterature (9). For caronaldehyde, the NMR spectrum wasmore complicated due to impurities, and a purity of 65%(hydrogen) was observed. It was later found that this samplecontained methanol and dichloromethane from the synthesis.In GC-MS, the peak area of caronaldehyde was typically 90%

of the total peak area in the chromatogram, and the MSspectrum was identical to the one published by Hakola et al.(9). Considering that a different spectral resolution was usedin the present case, the FTIR spectra for pinonaldehyde andcaronaldehyde match well those previously published byHakola et al. (9). The melting point of pinonaldehyde wasdetermined to be 250 ( 2 K.

Results and DiscussionTable 1 shows a summary of the conditions used in the variousexperiments.

Vapor Pressure. Vapour pressure measurements em-ploying the Knudsen effusion method are subjected to severalpossible systematic errors. One such error is connected tothe effective orifice size. The orifice is in principle a tubesince the foil is not infinitely thin. This creates a reduced fluxthrough the orifice not accounted for by the simple gas kinetictreatment. The fraction of the molecules predicted to reachthe orifice that is actually transmitted to the outside, the so-called Clausing factor W, can be estimated using eq II (18)where l is the foil thickness and d is the orifice diameter.

Applying eq II, W is 0.86 for the 0.3-mm orifice and 0.90 forthe 0.9-mm orifice. Equation II is approximate and based ontwo measured quantities. In addition, it is difficult to avoida slight deformation of the foil when making the orifice,thereby compromising its ideal geometry. It was thereforedecided to determine the effective orifice dimension bycalibration. High-purity naphthalene was used as thecalibration substance, and the measured vapor pressures,over a wide temperature interval, were compared to those byAmbrose et al. (19). The result is illustrated by Figure 2. Itwas found that the measured vapor pressures had to beincreased by a factor of 1.30 ( 0.10 and 1.10 ( 0.08 for the0.3- and 0.9-mm orifice, respectively, in order to match theliterature values. Expressed as Clausing factors, this corre-sponds to W ) 0.77 and W ) 0.91 for the two orifices andillustrates the value of the calibration procedure. The vaporpressures measured for the dicarbonyls were corrected byapplying these experimentally determined factors.

The loss of substance through the effusion orifice of theKnudsen cell creates a deviation from saturation pressureinside the cell. The saturation pressure is the quantity to bedetermined by the measurements, but it is the actual pressureinside the cell that is measured experimentally. One way ofinvestigating whether the difference between actual pressureand the saturation pressure is significant is to change theratio between the surface of the evaporating substance andthe orifice size.

Most of the measurements in the dicarbonyl case weremade by using the 0.3-mm orifice. In a few cases, the ratio

TABLE 1. Conditions Used in the Experimentsa

technique

relative rate

parametervapor

pressureIR

calibrationUV

absorption OH NO3

temperature (K) 255-276 300( 1 300 ( 1 298 ( 2 298 ( 2pressure (hPa) low <0.1 <0.1 1013 ( 5 1013 ( 5[pinonaldehyde]0 neat 0-6.8 2.4-14 8 9[caronaldehyde]0 neat 0-7.7 4.4-7.1 10-11 7-9[propene]0 10 10[CH3ONO] 40-52[NO] 2[N2O5] 40-45

a Concentrations are given in units of molecules cm-3 × 10-13.

ln[dicarbonyl]0

[dicarbonyl]t- t × kw )

k1

kpropene× ln

[propene]0

[propene]t(I)

CH3ONO + hν f CH3O + NO (4)

CH3O + O2 f HCHO + HO2 (5)

HO2 + NO f OH + NO2 (6)

W ) (1 + ld)-1

(II)

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was changed by a factor of 9 by using the 0.9-mm orificewithout any noticeable deviation of the measured vaporpressures from the best fit to the 0.3-mm data. In thenaphthalene case, the ratio was altered by changing theavailable surface of the compound. Normally, the naph-thalene was melted and allowed to solidify in a thin layer atthe bottom of the container to accomplish good thermalcontact. A few experiments were made by swirling the liquidduring solidification and allowing it to solidify, also on thewalls of the container, thereby increasing the evaporatingsurface by a factor of 4. No systematic difference was observedbetween these and the normal experiments. It is concludedthat the deviation from equilibrium pressure inside the cellwas not large enough to be significant neither for thedicarbonyl nor for the naphthalene measurements, and itfollows that the measured values are good estimates of thesaturation pressures.

Another cause of systematic errors is the use of impuresubstances. Three cases may be distinguished that affectsolid and liquid samples differently. An impurity may havea vapor pressure considerably higher, similar to, or consider-ably lower than the main constituent. Due to the lowdiffusional transport rates in solids, the first two cases areexpected to introduce errors of a magnitude equal to themolar ratio of the impurity multiplied by the difference inmolecular weights between main constituent and impurity.This is so because the availability of the impurity at the surfacewould be controlled by the loss rate of the main constituent.Only in the third case would large errors be expected sincehere the low volatility impurity would accumulate at thesurface, thereby affecting the surface composition. The maincomponent vapor pressure would then be seriously under-estimated. There were no indication of an impurity problemin the naphthalene experiments, and no impurity was detectedin the naphthalene by GC-MS analysis.

For a liquid sample and using the present equipment,diffusional mass transport takes place on a time scale shorterthan or similar to that of the surface loss to the gas phase.Here, the first case would give a rapid weight loss, corre-sponding to the mass of the high volatility impurity in thesample. This was observed in our experiments for bothdicarbonyls. The impurities were methanol and dichlo-romethane, both used as solvents in the synthesis but notcompletely removed in the workup of the preparations. Forpinonaldehyde, 3% of the initial charge was lost as high-volatility impurities while for caronaldehyde 15% was lost.After this rapid weight decrease, the loss rates stabilized, andsubsequent data were used to calculate vapor pressures. Inthe case of similar vapor pressures, both compounds wouldbe lost at a rate comparable to their molar ratios, and the

arguments given for solid samples would apply. Low-volatilityimpurities would accumulate in the sample and affect themeasurement if a large fraction of the sample was evaporated.The result would be a reduction of the measured vaporpressure with time.

Gas chromatographic analyses of the samples before andafter the Knudsen experiments showed the presence of onemain impurity in each dicarbonyl. Each impurity eluted after,but within 12 °C of the dicarbonyl in both cases. The relativeintensity of these peaks as compared to the main substancepeak were less than 4% and 7% for pinonaldehyde andcaronaldehyde, respectively. The impurity peak-main peakratio, measured before and after the Knudsen experiments,changed by less than 15% in both cases. On the basis ofthese observations, it is believed that the differences in vaporpressure between the dicarbonyl and its main impurity weresmall.

The vapor pressures of the two dicarbonyls as a functionof temperature were, as expected, found to follow theClausius-Clapeyron equation. The data are displayed inFigure 3. The coefficients in the equation were obtained byleast squares fitting and are given below (p in Pa, T in K; givenerrors are statistical at the 95% confidence level). Thesystematic errors are estimated to be less than 25% of themeasured values.

The heats of vaporization determined were 75.5 ( 5.6 and77.4 ( 6.9 kJ mol-1 for pinonaldehyde and caronaldehyde,respectively. Table 2 gives the saturation pressures over arange of temperatures. To the best of the authors knowledge,no vapor pressure data for 10-atom, dicarbonyl compoundsexist. The vapor pressure of decanal and decan-2-one at 283K were estimated to 6.7 and 13.6 Pa from data in Weast (20).As expected, the values for pinonaldehyde and caronaldehydeof 1 and 0.6 Pa at this temperature are lower than those forthe monocarbonyl compounds.

IR Calibration. Infrared spectra of pinonaldehyde andcaronaldehyde are shown in Figure 4. panels a and b. TheCO2 band at ∼2350 cm-1 is due to slight variations of CO2

concentration in the external transfer optics system. Pre-sented in Figure 5 are the absorbances at 1725.4 cm-1 forpinonaldehyde and at 1741.5 cm-1 for caronaldehyde vsconcentration multiplied by optical path length. Each datapoint in Figure 5 corresponds to one independent absorbance- pressure measurement. As is evident from Figure 5, the

FIGURE 2. Plot of ln (vapor pressure) vs temperature-1 fornaphthalene. The solid line shows a fit to the values of Ambroseet al. (19). Crosses are data from experiments with the 0.3-mm effusionorifice, and circles are data obtained with the 0.9-mm orifice.Pressures are given in Pa.

FIGURE 3. Plot of ln (vapor pressure) vs temperature-1 for thedicarbonyls. Crosses are data for pinonaldehyde, and circles aredata for caronaldehyde. Pressures are given in Pa.

pinonaldehyde

p ) exp(-9080((672)/T + 32.1((2.5)) (III)

caronaldehyde

p ) exp(-9315((835)/T + 32.3((3.1)) (IV)

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data yield straight lines with close to zero intercepts.Absorption cross sections found at 300 K are (in units of cm2

molecule-1, base e) (8.08 ( 0.32) × 10-19 at 1725.4 cm-1 and(3.29 ( 0.18) × 10-19 at 2714.4 cm-1 for pinonaldehyde and(1.04 ( 0.05)× 10-18 at 1741.5 cm-1 and (1.56 ( 0.09)× 10-19

at 2712.1 cm-1 for caronaldehyde. The given errors arestatistical at the 95% confidence level.

FTIR spectra of pinonaldehyde and caronaldehyde re-corded at atmospheric pressure were compared to spectraobtained at low pressure. No difference was observedbetween high- and low-pressure spectra. Hence, spectralsubtractions using either low-pressure or atmospheric pres-sure reference spectra give identical results. This is consistentwith infrared absorptions of large molecules not showing anypressure dependence, and it is concluded that the infraredcross sections are valid also at atmospheric pressure. Theinfrared cross sections were needed for the UV measurementsin this work but may also be used for in-situ determinationsof the gas-phase yield of pinonaldehyde or caronaldehyde ininvestigations of atmospheric oxidation of R-pinene and ∆3-

carene. The formyl C-H stretching band of pinonaldehydeat 2714.4 cm-1 shows unexpectedly strong absorption. Themeasured absorbtivity of this band is 1.67 times higher thanthe value proposed by Hatakeyama et al. (8), based on typicalabsorbtivities of other aldehydes.

UV Absorption Spectra. UV absorption spectra between275 and 350 nm of pinonaldehyde and caronaldehyde areshown in Figure 6. The shapes of the spectra are similar tothose of other aldehyde and ketone compounds (21, 22). Theabsorption maxima are located at 278 and 287 nm, which isin the range of other carbonyl compounds. Listed in Table3 are 5-nm interval averages of the absorption cross sections.Comparing the absorption cross section maxima from Figure6 with literature values of propionealdehyde (σ285 ) 61 ×10-21 cm2 molecule-1), acetaldehyde (σ290 ) 59 × 10-21 cm2

molecule-1), acetone (σ274 ) 52× 10-21 cm2 molecule-1), andmethyl ethyl ketone (σ276 ) 60 × 10-21 cm2 molecule-1) (asadapted from ref 22) strongly suggests that, as a firstapproximation, carbonyl group absorption is additive, i.e., acompound with two carbonyl groups has twice the absorptionof a single carbonyl group compound.

Relative Rate Experiments. Shown in Figure 7 are theplots of ln([caronaldehyde]0/[caronaldehyde]t) - tkw vs ln-([propene]0/[propene]t) for both OH and NO3 relative rateexperiments. The rate coefficients obtained from the relativerate experiments are presented in Table 4. The error limitsfor relative rate values are the 95% confidence interval fromthe least squares fitting. Given for comparison in Table 4 areliterature values (23) and estimates from a structure-reactivityrelationship (24).

TABLE 2. Vapor Pressure of Pinonaldehyde and Caronaldehydeas a Function of Temperaturea

temperature (°C) pinonaldehyde p (Pa) caronaldehyde p (Pa)

-10 0.09 0.05-5 0.2 0.1

0 0.3 0.25 0.6 0.3

10 1.0 0.615 1.8 1.020 3.0 1.825 5.1 3.030 8.3 5.035 13.6 8.2

a The measurements were mainly made between -18 and +3 °C.Thus, tabulated values outside this range have been extrapolated usingeqs III and IV.

FIGURE 4. (a) IR absorption spectrum of pinonaldehyde at 300 K and1013 hPa, obtained using a spectral resolution of 1 cm-1. (b) IRabsorption spectrum of caronaldehyde at 300 K and 1013 hPa, obtainedusing a spectral resolution of 1 cm-1.

FIGURE 5. IR calibration curves obtained at low pressure and at theindicated wave number by measuring absorbance as a function ofconcentration. Crosses are data for pinonaldehyde (1725.4 cm-1);circles are data for caronaldehyde (1741.5 cm-1).

FIGURE 6. UV absorption spectra. The upper curve is the averageof nine individual spectra of pinonaldehyde. The lower curve is theaverage of four individual spectra of caronaldehyde. The error barsshow the statistical error at the 95% confidence level. In order notto obscure the plot, error bars are only shown for each 5 nm.

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The measured values for the OH reactions are about afactor of 5 larger than calculated from the structure-reactivityrelation. The reason for this considerable deviation of theestimated value from the measured ones is not obvious. Animportant factor could be the possibility of formation ofcomparatively stable transition states in large molecules suchas pinonaldehyde and caronaldehyde. The formation of suchentities would increase the reaction rate. Also the ratecoefficients for the reaction with NO3 radicals are significantlyhigher than for other aldehydes (6, 23, 25), something thatalso could be explained by the argument given above.

For both OH and NO3 reaction with the dicarbonyls, themain product will originate from abstraction of the aldehydehydrogen atom. The hydrogen abstraction will, in air, leadto the formation of an acylperoxy radical. This radical is

expected to go through a series of reactions, splitting off amolecule of CO2 and eventually forming a new aldehyde. Ifenough NO2 is present, it is also possible to form a peroxynitrate in analogue with the formation of PAN (10, 26). ForOH, a minor route will be abstraction of the tertiary hydrogens(24).

Atmospheric Implications. The saturation pressures ofthe products are considerably higher than the productpressures anticipated from oxidation reactions in the atmo-sphere. The original terpene is not often present at pressuresabove 10-3 Pa (0.01 ppm), and the product concentrationsare not expected to exceed those of the starting material byseveral orders of magnitude. The vapor pressure data showthat the main products from atmospheric oxidation ofR-pinene and ∆3-carene can neither nucleate homogeneouslynor condense on other nuclei. Other mechanisms are thusneeded to explain their presence in particulate matter (3).

Raoult’s law was applied to partitioning of organiccompounds between gas and condensed “particulate” phaseby Pankow (27, 28) and further developed by Odum et al.(29). Pankow (28) defined an equilibrium partitioningconstant Kp,i as

where Fi,om is the mass concentration of the condensedcompound i, TSP is the total suspended particulate concen-tration having a fraction fom of organic matter (om), and Ai

is the mass concentration of compound i in the gas phase.Further, MWom is the average molecular weight of thecondensed organic matter, úi is the activity coefficient ofcompound i in the organic matrix, p°i is the vapor pressureover pure i at temperature T, and R is the common gasconstant. Expression V confirms the intuitive concept thatsubstances with a high vapor pressure should be present inlow concentrations in the condensed phase unless the activitycoefficient for i in the actual combination of substances ismuch smaller than 1. If it is assumed that MWom ) 0.2 kgmol-1, which is reasonable unless polymerization occurs, andfom ) ú )1, then Kp,pinonaldehyde ) 1.2 × 104 m3 kg-1 andKp,caronaldehyde ) 2.0 × 104 m3 kg-1 at 283 K.

Expression V may be applied to data given by Yokouchiand Ambe (3) concerning pinonaldehyde found in aerosolparticles. These authors observed 2 ng m-3 particulatepinonaldehyde, constituting about 0.5% of the aerosol organicmatter. This corresponds to a TSP concentration of about0.4 µg m-3 with fom ) 1. At equilibrium, this would requirea gas-phase concentration of 4.2×10-7 kg m-3, correspondingto about 60 ppb. The gas-phase concentration needed tomaintain the observed condensed-phase concentration ap-pears to be unrealistically high unless the activity coefficientis considerably less than 1.

Odum et al. (29) obtained effective equilibrium constantsfrom fitting a two-product model to experimental aerosoldata for R-pinene. Their data are well described by equi-librium constants of 1.71 × 108 and 4 × 1010 m3 kg-1 at 308K. Employing expression V and the assumptions made aboveabout MWom, fom and ú, this results in vapor pressures of 7.5

TABLE 3. UV Absorption Cross Sections (Base e)a

wavelengthinterval (nm)

pinonaldehydeσ × 1019

(cm2 molecule-1)

caronaldehydeσ × 1019

(cm2 molecule-1)

275-280 1.219 ( 0.029 1.240 ( 0.065280-285 1.296 ( 0.023 1.200 ( 0.051285-290 1.348 ( 0.044 1.168 ( 0.037290-295 1.294 ( 0.057 1.083 ( 0.042295-300 1.217 ( 0.048 0.977 ( 0.041300-305 1.062 ( 0.041 0.822 ( 0.039305-310 0.894 ( 0.034 0.650 ( 0.041310-315 0.684 ( 0.026 0.469 ( 0.037315-320 0.490 ( 0.020 0.316 ( 0.033320-325 0.307 ( 0.015 0.165 ( 0.032325-330 0.172 ( 0.014 0.072 ( 0.029330-335 0.072 ( 0.011335-340 0.029 ( 0.008340-345 0.008 ( 0.005

a The cross sections are the average over a 5-nm interval. The errorsare statistical, at the 95% confidence level, from the average of thedifferent spectra measured (9 for pinonaldehyde and 4 forcaronaldehyde).

FIGURE 7. Plot of ln ([caronaldehyde]0/[caronaldehyde]t) - tkw vs([propene]0/[propene]t) in the relative rate experiments. Circles aredata for the reaction with hydroxyl radicals; crosses are data forthe reaction with nitrate radicals.

TABLE 4. Rate Coefficient Obtained in Relative Rate Experimentsa

reaction this work Glasius et al. (23)estd using structure-reactivity

relationship

pinonaldehyde + OH (8.72 ( 1.14) × 10-11 (9.1 ( 1.8) × 10-11 2.1 × 10-11

caronaldehyde + OH (1.21 ( 0.36) × 10-10 2.1 × 10-11

pinonaldehyde + NO3 (2.35 ( 0.37) × 10-14 (5.4 ( 1.8) × 10-14

caronaldehyde + NO3 (2.71 ( 0.15) × 10-14

a The rate coefficients are given in cm3 molecule-1 s-1. The errors are at the statistical 95% confidence level from the least squares fitting of therelative rate plots. Shown for comparison are the available literature values and estimations using a structure-reactivity relationship (24).

Kp,i )Fi,om

AiTSP)

RTfom

MWomúip° (V)

VOL. 31, NO. 11, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 3171

× 10-5 and 3.2 × 10-3 Pa. Comparing these pressures withthat for pinonaldehyde determined in this work, it appearsthat pinonaldehyde is not contributing to the observedaerosol.

Both compounds show a tendency for adsorption to solids,illustrated by the wall loss in the chamber experiments, andone conceivable physical removal processes may be drydeposition.

The quantum yields for dissociation were not determinedin this work. However, if a yield of unity is assumed,atmospheric lifetimes with respect to photolysis can becalculated using actinic fluxes from the literature (30). Theselifetimes will be the lower limits since a quantum yield of 1is most likely an overestimation. For small aldehydes, theyields for dissociation are significantly smaller than 1, whereaslittle is known for aldehydes of molecular weight comparableto present compounds (31). The lower limit-photolyticlifetime ()kphotolys

-1 ) for a range of conditions are given inTable 5. The large difference in lifetime (≈50%) betweenpinonaldehyde and caronaldehyde is caused by the differencein absorption cross section in the high wavelength region(i.e., 320-350 nm).

Table 5 indicates that the dominant gas-phase removalprocesses in the atmosphere, both for pinonaldehyde andcaronaldehyde, will be reaction with OH radicals and possiblyalso photodissociation. The natural lifetimes are estimatedto be of the order of several hours during daytime. At nightthey will disappear slowly, with natural lifetimes of 0.5 dayby reaction with NO3 radicals. It must be emphasized thatthe radical concentrations, needed for the calculation of thelifetimes regarding OH and NO3 reactions, are dependent onatmospheric conditions and that the concentrations may varysubstantially. In Table 5, the concentrations used arerepresentative for day time (OH) (22) and nighttime (NO3) (6)local peak values at mid-latitudes. The expected stableorganic products, both from photolysis and radical attack,are new carbonyl compounds with one carbon atom less,i.e., with a slightly lower molecular weight than the originaldicarbonyls. Such compounds should be volatile enough tostay in the gas phase. They may therefore well be availablefor participation in further peroxy radical formation and thuscontribute to tropospheric ozone formation.

AcknowledgmentsThis work was supported financially through the EU Envi-ronment and Climate BIOVOC project. I.W. thanks the

European Commission for a Human Resources and Mobilitygrant. Thanks are due to the research group of Prof. K. H.Becker, Department of Physical Chemistry, University ofWuppertal, for permission to use their UV absorption setup.In particular, the help and advice by Dr. H. G. Libuda and Dr.I. Barnes are gratefully acknowledged.

Literature Cited(1) Guenther, A.; Hewitt, C. H.; Erickson, D.; Fall, R.; Geron, C.;

Graedel, T.; Harley, P.; Klinger, L.; Lerdau, M.; McKay, W. A.;Pierce, T.; Scholes, B.; Steinbrecher, R.; Tallamraju, R.; Taylor,J.; Zimmerman, P. J. Geophys. Res. 1995, 100, 8873.

(2) Atkinson, R. J. Phys. Chem. Ref. Data 1994, Monograph 2.(3) Yokouchi, Y.; Ambe Y. Atmos. Environ. 1985, 19, 1271.(4) Warnek, P. Heterogeneous and Liquid-Phase Processes;

Springer-Verlag: Berlin-Heidelberg, 1996.(5) Hatakeyama, S.; Izumi, K.; Fukuyama, T.; Akimoto, H. J. Geophys

Res. 1989, 94, 13013.(6) Wayne, R. P.; Barnes, I.; Biggs, P.; Burrows, J. P.; Canosa-Mas,

C. E.; Hjorth, J.; LeBras, G.; Moortgat, G. K.; Perner, D.; Poulet,G.; Restelli, G.; Sidebottom, H. Atmos. Environ. 1991, 25A, 1.

(7) Arey, J.; Atkinson, R.; Aschmann, S. M. J. Geophys Res. 1990, 95,18539-18546.

(8) Hatakeyama, S.; Izumi, K.; Fukuyama, T.; Akimoto, H.; Washida,N. J. Geophys Res. 1991, 96, 947.

(9) Hakola, H.; Arey, J.; Aschmann, S. M.; Atkinson, R. J. Atmos.Chem. 1994, 18, 75.

(10) Wangberg, I.; Barnes, I.; Becker, K. H. Environ. Sci. Technol.1997, 31, 2130-2135.

(11) Jansson, R. J. Geophys. Res. 1993, 98, 2839.(12) Libuda, H. G.; Sabel, F.; Fink, E. H.; Becker, K. H. J. Phys. Chem.

1990, 94, 5860.(13) Wangberg I.; Ljungstrom E.; Olsson, B. E. R.; Davidson, J. J. Phys.

Chem. 1992, 96, 7640.(14) Langer, S.; Ljungstrom, E. J. Phys. Chem. 1994, 98, 5906.(15) Canosa-Mas, C. E.; Smith, S. J.; Waygood, S. J.; Wayne, R. P. J.

Chem. Soc. Faraday Trans. 1991, 87, 3473.(16) Black, A. P.; Babers, F. H. Organic Synhesis; Johnson, J. R., Ed.;

Wiley: New York, 1939; Vol. XIX, p 64.(17) McMurry, J. E.; Bosch, J. Org. Chem. 1987, 52, 4885.(18) Clausing, P. Ann. Physik, 1932, 12, 961.(19) Ambrose, D.; Lawrenson, I. J.; Sprake, C. H. S, J. Chem.

Thermodyn. 1975, 7, 1173.(20) Weast, R. C., Ed. Handbook of Chemistry and Physics, 52nd ed.;

The Chemical Rubber Company: Cleveland, 1972; p D166.(21) Lee, E. K. C.; Lewis, R. S. Adv. Photochem. 1980, 12, 1.(22) Finlayson-Pitts, B. J.; Pitts Jr., J. N. Atmospheric Chemistry;

Wiley: New York, 1986.(23) Glasius, M.; Calogirou, A.; Jensen, N. R.; Hjorth, J.; Nielsen, C.

J. Int. J. Chem. Kinet. 1997, 29, 527.(24) Kwok, E. S. C.; Atkinson, R. Atmos. Environ. 1995, 29, 1685.(25) Atkinson, R. J. Phys. Chem. Ref. Data 1991, 20, 459.(26) Stephens, E. R. Adv. Environ. Sci. 1969, 1, 119.(27) Pankow, J. F. Atmos.Environ. 1994, 28A, 185.(28) Pankow, J. F. Atmos.Environ. 1994, 28A, 189.(29) Odum, J. R.; Hoffmann, T.; Bowman, F.; Collins, D.; Flagan, R.

C.; Seinfeld, J. H. Environ. Sci. Technol. 1995, 30, 2580.(30) Demerjian, K. L.; Schere, K. L.; Peterson, J. T. Adv. Environ. Sci.

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Received for review February 21, 1997. Revised manuscriptreceived July 1, 1997. Accepted July 23, 1997.X

ES970151A

X Abstract published in Advance ACS Abstracts, September 15, 1997.

TABLE 5. Estimated Lifetimes for Pinonaldehyde andCaronaldehyde with Respect to Photodissociation, Reactionwith OH, and Reaction with NO3

a

photodissociationreaction with

OH [OH]) 1 × 106

molecules cm3

reaction withNO3 [NO3]) 1 × 109

molecules cm3Jul 1,50° N

Jul 1,60° N

Jan 1,50° N

pinonaldehyde 3.2 h 11 h 3.3 h 3.9 h 22 hcaronaldehyde 2.3 h 10 h 5.8 h 6.8 h 43 h

a The lifetimes for reaction with OH and NO3 were calculated usinglifetime ) (koxidant × [oxidant])-1. The photodissociation lifetimes werecalculated using the absorption cross sections in Table 3 and actinicfluxes from Demerjian et al. (30) at noon for locations and dates specified.A quantum yield of 1 for photodissociation was used.

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