Environ. Scl. Technol. 1983, 17, 50-51
Rates of Reaction of Hydroxyl Radlcals with 2-(Dimethylamino)ethanol and 2-Amino-2-methyl-1-propanol in the Gas Phase at 300 f 2 K Geoffrey W. Harrls and James
N. Pltts, Jr."
Statewide Air Pollution Research Center, University of California, Riverside, California 92521 The rates of the reactions of OH radicals with 2-(dimethy1amino)ethanol and 2-amino-2-methyl-1-propanol have been determined by using a flash photolysis-resonance fluorescence technique to be (4.7 f 1.2) X lo-" and (2.8 f 0.5) X cm3 molecule-' s-', respectively, at 300 f 2 K. Implications of the increased use of these compounds (as paint solubilizers) on air quality are discussed. The recent proposal by the California Air Resources Board (1)of a new policy to control solvent emissions from assembly line paint operations has led to the increased use of low-solvent (i.e., waterborne and high solids) coatings for these purposes. Such coatings contain (as solubilizers) significant quantities of amines, in particular 2-(dimethylamino)ethanol, (CH3)2NCH2CH20H(DMAE), and to a lesser extent, its structural isomer 2-amino-2methyl-1-propanol, (CH3)2C(NH2)CH20H (AMP). The occurrence of such amines in these finishes raises the possibility that one air quality effect (i.e., reduction of oxidant formation) may be traded for another potential health effect (i.e., formation of nitrosamines or nitramines) from amine precursors and oxides of nitrogen in air (2,3). The major loss process for both DMAE and AMP in polluted urban atmospheres is expected to be reaction with the hydroxyl radical. Thus, we report here results of direct measurements of the absolute rate constants for the reaction of DMAE and AMP with hydroxyl radicals at 300 f 2 K using the flash photolysis-resonance fluorescence technique. The experimental technique employed involves formation of OH radicals in a flow system by pulsed vacuum ultraviolet photolysis of H20 at X 2105 nm and monitoring their resulting decay rates by detection of the OH resonance fluorescence band at 306.4 nm (A28+,J = 0 X211, J' = 0) with a cooled photomultiplier tube. This technique has been described in detail previously (5-7) and thus need not be discussed further here. The reaction was studied at 50 f 1torr total pressure of argon. The concentration of the amines in the argon flow was determined by UV absorption at X 200-225 nm. The observed decays of the hydroxyl radical concentration were always exponential, and the first-order decay rates are shown as a function of DMAE and AMP concentration in Figures 1 and 2, respectively. Variation of the flash energy by a factor of 2 had no effect on the decay rates within experimental error nor had an increase in the total pressure of argon to -100 torr. The slopes of the least-squares lines shown in Figures 1 and 2 yield the following rate constants for the reaction of hydroxyl radicals with DMAE and AMP at 300 f 2 K: kl(DMAE) = (4.7 f 1.2) X lo-'' cm3 molecule-' s-l
-
kl(AMP) = (2.8 f 0.5)
X
lo-''
cm3 molecule-' s-'
The error limits reflect twice the standard deviation of the slopes in Figures 1and 2 plus the estimated uncertainties (10%) involved in measuring flow rates, pressures, and optical density at the wavelengths used. This is the fiist rate measurement for the reaction of OH with AMP to be reported in the literature, but Anderson 50
Environ. Sci. Technol., Vol. 17, No. 1, 1983
and Stephens (8) report a value of ( 8 3 X cm3 molecule-' s-' for OH + DMAE over the temperature range 269-364 K, which is about 40% higher than that obtained in this study. The reason for this difference is not clear, especially since the experimental method (8) was very similar to the present one (5-7). Taking the average hydroxyl concentration in the urban atmosphere to be -3 X lo6 molecules cm-3 leads to estimated lifetimes (l/e) for DMAE and AMP of approximately 2 and 4 h, respectively, under such conditions, for oxidation by OH radicals. It is interesting to speculate on the fate of the radicals produced in these reactions under atmospheric conditions. Although no mechanistic information was obtained in this study, comparison of the measured rate constants with those for OH reactions of other molecules containing similar groups can provide insight into the branching ratios for the possible mechanisms. Thus, consideration of the rate for OH + trimethylamine cm3molecule-' s-' (9))leads to the conclusion (6.2 X that H abstraction from each N-methyl group contributes -2.0 X cm3 molecule-' s-' to the overall rate (or -0.67 X lo-'' cm3molecule-' s-' per C-H bond). Further, the reactivity of the alcohols methanol, ethanol, and propanol toward OH (10, 11)may be used to deduce that the -CH2CH20H group contributes -0.3-0.4 X 10-l' cm3 molecule-l s-l to the reactivity of n-propanol. Hence, we may estimate the rate constant for OH + DMAE to be -6(0.67 X lo-'') + (0.4 X 10-11) 4.4 X lo-'' cm3molecule-l s-', in reasonable agreement with the measured value. The slightly higher value of the measured rate constant may reflect higher reactivity for the N-bonded rather than C-bonded -CH2CH20H group. On the basis of this argument, we may conclude that H abstraction from the N-methyl groups dominates in the reaction of OH + DMAE and constitutes -60430% of the overall reaction. The subsequent reactions of the radical produced in this dominant path may be as follows:
-
I
I1 CH3kHzCH20H
t
HCHO
I11 where scission of radical I1 is expected to occur more rapidly than H abstraction by O2from the radical center. The atmospheric chemistry of the dimethylamino (DMA) radical, which is analagous to the N-centered radical 111, has been investigated by previous workers (2,3,12), who have concluded (12)that addition of NO and NOz to DMA to form nitrosamines and nitramines is in excess of lo7 times faster than reaction with 02.In sunlit atmospheres, which are a prerequisite for significant OH concentrations, the nitrosamine formed from I11 will rapidly photodissociate, thus it is likely that the major N-containing product from DMAE will be the photostable nitramine derivative of 111, CH3(CH2CHzOH)NNOZ. Arguments based on the OH rate constants for methylamine (2.2 X 10-l' cm3molecule-l s-l (13)),ethylamine (2.8
0013-936X/83/09 17-0050$01.50/0
0 1982 American Chemical Society
500
r
T.300 f 2K
OH + DMAE
O/
/--
400
IV
IV
Figure 1. Psuedo-first-order decay rates of the hydroxyl radical concentration as a function of DMAE concentration at 300 f 2 K. I
CH+AMP
-CH3
-CHzOH
CH3
'c=N-H /
CH3\ ,C=N-H
(3)
c A3
as aid in assessing the atmospheric yield of the nitramine from IV (analogues of which are carcinogenic in laboratory animals (14, 15)). Acknowledgments We thank R. Atkinson, W. P. L. Carter, and A. M. Winer for assistance in preparation of this manuscript.
T=300iZK
Registry No. 2-(Dimethylamino)ethanol,108-01-0; 2-amino2-methyl-7-propano1, 124-68-5; OH., 3352-57-6.
Literature Cited
,
Calif. Air Resour. Board Bull. 1978, 9, 1. Pitts, J. N., Jr.; Grosjean, D.; Van Cauwenberghe, K.; Schmid, J. P.; Fitz, D. R. Environ. Sei. Technol. 1978,12, 946.
Tuazon, E. C.; Winer, A. M.; Graham, R. A.; Schmid, J. P.; Pitts, J. N., Jr. Environ. Sei. Technol. 1978, 12, 954. Finlayson, B. J.; Pitts, J. N., Jr. Science (Washington,D.C.)
/
d O'
1
d
k
[AMP]/molec
L
1976,192, 111. Ib
;1
cm-3
Figure 2. Pseudo-firstorder decay rates of the hydroxyl radical concentration as a function of AMP concentrations at 300 f 2 K.
x
cm3molecule-1 s-l (11))and for ethanol (0.33 X cm3 molecule-l s-l (10, 11)) indicate that, although abstraction from C-H bonds usually is faster than from N-H bonds in primary amines, in the case of AMP the absence of hydrogen on the a carbon suggests that C-H abstraction would also be slow. Thus the reaction would most likely proceed by abstraction from the primary amino group giving radical IV. CH3
I .
CH3-C-NH
I CH20H
IV
The only likely fates of radical IV under atmospheric conditions seem to be given in eq 1-3. The suggested pathways involving loss of methyl (or methanolic) groups are somewhat unusual, and product studies of the photooxidation of AMP (or of the more symmetrical tert-butylamine) may prove instructive as well
Atkinson, R.; Perry, R. A.; Pitts, J. N., Jr. J . Chem. Phys. 1976, 65, 306. Atkinson, R.; Hansen, D. A.; Pitts, J. N., Jr. J. Chem.Phys. 1975, 62, 3284.
Atkinson, R.; Hansen, D. A,; Pitts, J. N., Jr. J. Chem. Phys. 1975, 63, 1703.
Anderson, L. G.; Stephens, R. D. 14th Informal Conference on Photochemistry, Newport Beach, CA, Apr, 1980. Atkinson, R.; Perry, R. A.; Pitts, J. N., Jr. J. Chem. Phys. 1978, 68, 1850.
Campbell, I. M.; McLaughlin, D. F.; Handy, B. J. Chem. Phys. Lett. 1978, 53, 385. Overend, R.; Paraskevopoulos, G. J. Phys. Chem. 1978,82, 1329.
Lindley, C. R. C.; Calvert, J. G.; Shaw, J. H. Chem. Phys. Lett. 1979, 67, 57. Atkinson, R.; Perry, R. A.; Pitts, J. N., Jr. J. Phys. Chem. 1977, 66, 1578. Goodall, C. M.; Kennedy, T. H. Cancer Lett. 1976,1,296. Druckrey, H.; Preussmann, R.; Schmahl, D.; Muller, M. Naturwissenschaften 1961,48, 134. Received for review May 14,1982. Accepted August 30,1982. This work was supported in part under California Air Resources Board Research Contract A7-175-30 and in part under National Science Foundation Grant CHE 79-10447. Mention of brand names or products does not imply endorsement or recommendation for use by either agency.
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