(8) Dams, R., Rahn, K. A., Winchester, J. W., Enuiron. Sci. Technol., 6,441-48 (1972). (9) Paciga, J. J., Chattopadhyay, A., Jervis, R. E., in “Nuclear Methods in Environmental .Research-11”, J. R. Vogt and W. Meyer, Eds., pp 286-300, University of Missouri, Columbia, Mo., 1974. (10) Hoffman, G. L., Duce, R. A,, Enuiron. Sci. Technol., 5 , 1134 (1971). (11) Gordon, G. E., Zoller, W. H., Gladney, E. S., in “Trace Substances in Environmental Health-VII”, D. D. Hemphill, Ed., pp 167-75, University of Missouri, Columbia, Mo., 1973. (12) John, W., Kaifer, R., Rahn, K., Wesolowski, J . J., Atmos. Enuiron., 7, 107-18 (1973). (13) Pattenden, N. J., presented a t the 41st Ann. Conf. of the National Society for Clean Air, 14-18 October, Cardiff, UK, 1974. (14) Bogen, J., Atmos. Enuiron., 7,1117-25 (1973). (16) Wedepohl, K. H., in “Origin and Distribution of the Elements”, L. H. Ahrens, Ed., pp 999-1016, Pergamon Press, Oxford, England, 1968.
(16) Lee, R. E., Jr., Goranson, S., Emiron. Sci. Technol., 6,1019-24 (1972). (17) Whitby, K. T., Charlson, R. E., Wilson, W. E., Stevens, R. K., Science, 183,1098-99 (1974). (18) Lee. R. E.. ibid.. t m 1099-100. (19) Martens, C. S., WAolowski, J. J., Kaifer, R., John, W., Enuiron. Sci. Technol.. 7,817-20 (1973). (20) Zoller, W. H., Gladney, E. S., Gordon, G. E., Bors, J. J., in “Trace Substances in Environmental Health-VIII”, D. D. Hemphill, Ed., pp 167-72, University of Missouri, Columbia, Mo., 1974. (21) Roberts, T. M., Hutchinson, T. C., Paciga, J. J., Chattopadhyay, A., Jervis, R. E., Van Loon, J., Parkinson, D. K., Science, 186, 1120-23 (1974);Institute for Environmental Studies Report E E # 1, University of Toronto, Toronto, Canada, 1975. (22) Paciga,-J.J.,Roberts, T. M., Jervis, R. E., Enuiron Sci. Technol., 9,1141-44 (1976). Received for review August 24, 1975. Accepted April 30, 1976
Reactions Which Relate to Environmental Mobility of Arsenic and Antimony. II. Oxidation of Trimethylarsine and Trimethylstibine George E. Parris” and F. E. Brinckman Inorganic Chemistry Section, National Bureau of Standards, Washington, D.C. 20234
w Oxidation of trimethylamine and trimethylstibine by atmospheric oxygen and other reagents is examined, and semiquantitative rate constants are calculated. In methanol solution the rate constant for reaction of dissolved oxygen with trimethylstibine is greater than lo-* M-’ s-l, whereas for the oxidation of trimethylamine, the rate constant is less than lo-* M-’ s-l. In the gas phase the rate constants are estimated as 10,’’and M-l s-l for reaction of trimethylstibine and trimethylarsine, respectively, with oxygen. A scheme based on PMR evidence for reactive intermediates is suggested to account for the products of oxidation, (CH3)3EO and (CH:J2E02H,of these compounds (CH&E, E = Sb, As. From these results, even if biological methylation of antimony occurs in nature analogous to that of arsenic, the rapidity with which (CH:I):$b is oxidized would probably prevent hazardous concentrations from building up in aerated surroundings.
in the chemical and physical properties of organic arsenic and antimony compounds because these properties will ultimately determine the pattern of environmental distribution of arsenic and antimony mobilized via biological methylation (Scheme 1). Scheme 1 Inorganic As a n d S b Pollutants
/ / I
\ \
Abiotic Reactions with Natural and Man-Made Chemicals
Heavy metals and metalloids such as arsenic and antimony have always been part of the environment. However, contemporary industrial utilization has redistributed these elements so that in many cases they now intermingle more extensively with biological food chains. For example, arsenicbased pigments were once used in buildings, arsenic-based pesticides are currently in use, and antimony oxide is widely used in flame retardant systems. It was originally believed that the chemical form in which the metals and metalloids were discharged as pollutants (inorganic oxides) rendered them inert, immobile, and relatively nontoxic. Unfortunately, in the case of many metals and metalloids including arsenic (and by inference chemically similar antimony), microorganisms can reduce and methylate the inorganic compounds to give organometallic species which are stable and mobile in water and air (1, 2 ) (at this time it has been shown that many heavy metals and metalloids including S, Se, Te, As, Sn, Pb, and Hg are subject to biological methylation). In addition, the physical properties of organometallic compounds often allow them to penetrate some of the barriers which higher organisms normally apply against toxic substances (the blood-brain barrier). Our interest is to examine reactions which provide models for environmental alterations 1128
Environmental Science & Technology
Salts,
Oxides,
Acids,
Complexes,
Ligands
In Part I of this series, quaternization of trimethylamine and trimethylstibine was studied ( 3 ) . In this part, abiotic oxidation will be examined. Abiotic oxidation of alkyl-arsines and -stibines can be achieved with a variety of reagents including atmospheric oxygen, hydrogen peroxide, and mercuric oxide. Oxidation by mercuric oxide ( 4 , 5 )is reported to give high yields of the arsine or stibine oxide (R3E0, E = As or Sb) with no cleavage of the arsenic-carbon or antimony-carbon bonds. Oxidation by hydrogen peroxide (6, 7) yields the arsine and stibine oxides, but cleavage of the As-C or Sb-C bonds occurs when excess peroxide is used and especially in the presence of base. Oxidation by atmospheric oxygen (6, 8, 9) yields a mixture of the arsine or stibine oxide and the arsinic or stibinic acid (R2E03H). Before presenting our results on these reactions, it is desirable to introduce some of the descriptive chemistry ( 1 0 ) of the products of the oxidation. There appears to be some double bond character in the As=O bond of trialkylarsine oxides (R3As+-O-
-
RaAs=O). These compounds have been used to form numerous transition metal complexes, and in the presence of acid, the oxides appear to be protonated. In aqueous solution, compounds of the type R3AsXz (X = incipient univalent anion) are rapidly hydrolyzed according to the sequence ( I I ) la-lc.
Thus, depending upon pH, these solutions can be regarded as solutions of R3AsO or R3AsX2. There is less tendency for antimony to form a double bond to oxygen. Morris et al. (5) have studied trimethylstibine oxide and found that, when prepared by oxidation of trimethylstibine with HgO in anhydrous diethyl ether, the compound has a dimeric structure. The compound is very hygroscopic, and in the presence of protic solvents (HOR where R = H, organogroup), the following rapid hydrolysis occurs.
These solutions can be regarded as solutions of the stibine oxide as long as equilibrium 2 is kept in mind. In the case of the arsinic and stibinic acids, the differences in the tendency of antimony and arsenic to form double bonds to oxygen are also manifest. Dimethylarsinic acid (cacodylic acid) is very soluble in water.
It is a hydrogen-bonded dimer in the solid phase (12),though it is monomeric in protic solvents. The pK of dimethylarsinic acid is about 6.2 a t zero ionic strength. The acid can be protonated to give (CH3)2As+(OH)zwhich has a pK of about 1.6 at low ionic strength ( I O ) . In contrast, dialkylstibinic acids are polymeric with low solubility in water and even lower solubility in organic solvents. Stibinic acids precipitated from aqueous solution have water associated with them due to formation of hydrated oligomers and monomers which lose water and polymerize when heated or dried.
Experimental Trimethylstibine and trimethylamine solutions in methanol were prepared and standardized by the techniques described in Part I ( 3 ) .Proton magnetic resonance spectra were made on commercial 60 MHz and FT 90 MHz instruments. In bench scale oxidations of trimethylstibine by air, the solutions were obtained by distilling the ether/(CHa)sSb mixture from the Grignard reaction mixture used to prepare trimethylstibine from SbClx+CH:jMgBr (13). In one case, an anhydrous solution of trimethylstibine was exposed to a stream of air dried by passage through concen-
trated sulfuric acid. In another case, a layer of 10% aqueous hydrochloric acid was added to the ether solution before exposing the mixture to ambient air. In both cases, the reactions were worked up by filtration of the solids and evaporation of the liquid phases to isolate products. In both cases, only traces of material were found in the ether layers. The aqueous layer from the protic oxidation system contained 1.4 g of (CH3)3SbC12. The solid isolated by filtration of the aqueous HCl/ether reaction mixture yielded 6.0 g of (CH3)3SbC12 (Theoretical % C 15.15, % H 3.82; Found % C 15.31,% H 3.85) after recrystallizations from methanol and acetone had separated 0.3 g of an insoluble white solid (Found % C 4.39, % H 1.84). Thus, oxidation under low pH protic conditions resulted in only a very small amount of cleavage of carbon-antimony bonds and produced mainly (CH3)3SbO isolated as the dichloride. In the anhydrous oxidation procedure, the solid isolated by filtration of the reaction mixture was treated with 10% aqueous hydrochloric acid and was recrystallized from methanol and acetone to yield as the only product 10.5 g of (CH3)3SbC12 (Found % C 15.23, % H 3.79). Thus, also under anhydrous conditions there is very little carbon-antimony bond cleavage during atmospheric oxidation of (CH3)3Sb. The small, NMR-scale oxidation studies were conducted by placing approximately 1.0 ml of a 0.5 M solution of trimethylstibine or trimethylamine in methanol in a thin-wall NMR tube inside an inert (N2) atmosphere box. The tube's inside diameter was 4 mm, the liquid depth was about 4 cm, and the surface of the liquid was about 12 cm from the opening of the tube. The tubes were removed from the drybox and uncapped in an efficient hood for timed periods of exposure to the atmosphere. Experiments involving gaseous reactants were accomplished with a conventional all glass high-vacuum line employing grease-free stopcocks. Pressure changes were measured with a telescopic tensimeter (f0.02 mm). Gas-phase oxidation of trimethylamine was studied by filling a reaction bulb attached to a mercury manometer with 158.5 torr of (CH3)3As gas, then condensing the (CH3)3As in a narrow section of the reaction bulb, and admitting 173.1 torr of oxygen. When the system was allowed to warm up, the total pressure was 333.6 torr at 24.5 "C. [The difference between the measured partial pressures and total pressure is due to the cooling effect of the condensed (CH3)3Ason part of the oxygen.] Calculated from the pressures, the initial concentrations M and 9.43 X M, of (CH&As and 0 2 were 8.54 X respectively. The system was kept between 25 and 21 "C in the dark 90% of the time and exposed to the laboratory lights about 10% of the time. The pressure (corrected for temperatorr/s over a ture) decreased very slowly a t a rate of 5 x six-day period. The equation
can be used to relate change in pressure to the changes in concentration of the gases. T I similar experiments with (CH&Sb, very rapid oxidations fot .)wed by ignitions were observed. In a competition experiment a mixture of roughly equal amounts of (CH3)$3b and (CH3)3Aswere condensed in a reaction bulb, and oxygen was added. As the mixture of (CH&Sb and (CH&As evaporated, formation of a white smoke was followed by ignition producing a black coating on the flask. The gases were analyzed qualitatively by infrared spectroscopy (50 torr, 5-cm path length, KBr windows). The following assignments were made by comparison with spectra of authentic gas samples: (CH:1):$b (2950, 2880, 1440,1220,1140,810,510cm-l) and (CH3):iAs (2950, 2870, 1420, 1270, 890,570 ern-') occurred in roughly equal amounts and were the dominant gases present. Methane Volume 10, Number 12, November 1976
1129
CH4 (3020,1290 cm-’), ethylene CzH4 (952 cm-I), and acetylene C2H2 (730 cm-l) were tentatively identified. Only a trace of carbon monoxide CO (2150 cm-l) and even less COP (2380 cm-l) were detected. No water, methanol, or ethane were detected. Extraction of the solids coating the reaction bulb with methanol and examination of the PMR spectrum revealed only (CH3)3AsO and (CH&SbO in a ratio of about 1:lO.
Results and Discussion Trimethylamine and trimethylstibine can be oxidized by a number of reagents as discussed in the introduction, but oxidation by atmospheric oxygen is the most important reaction in the environment. The technique used to follow oxidation of (CH3)sAs and (CH&Sb in solution was merely to put about 1ml of 0.5 M solution into a thin-walled NMR tube and follow the changes in the PMR spectrum as a function of “exposure time”, Le., the length of time the NMR tube was allowed to stand uncapped in an efficient hood. During the oxidation of (CH&Sb (0.74 6) in methanol, the following observations were made (Figure 1).In the very early stages of the reaction (exposure time < 10 min), signals for (CH:&$bO (1.63 6) and a species believed to be (CH&SbOR, where R = H or OCH3 (1.00 6), were observed with approximately the same intensities. Both signals increased during the early stages of the reaction, but the rate of increase of the (CH,J:$bO signal was much greater. In the middle part of the reaction (spectrum C, Figure l), interruption of the oxidation process for days had no effect upon the relative intensities of the signals. This result indicates that the signal at 1.00 6 is due to a species which results from reaction of (CH3)3Sb with oxygen and which requires additional oxygen for itself to react (Le., in the absence of oxygen it is indefinitely stable). In the early stages of the reaction, an insoluble white solid began to form, but only near the gadliquid interface. After the middle part of the reaction, solid formed more or less uniformly throughout the liquid phase. This result suggests that during the early part of the reaction, the rate was controlled by diffusion of oxygen into the liquid, and during the latter stages of the reaction, the rate was determined by actual kinetic factors in the liquid phase. The signal at 1.00 6 reached a maximum during the middle part of the reaction, but it did not decrease until the signal for (CHA);$b had decreased to a similar intensity (Figure 2). In the latter stages of the reaction, the signal for (CH3)3Sb and the signal at 1.00 6 decreased at about the same rate. Addition of fresh (CH&Sb to an exhaustively oxidized solution did not regenerate the 1.00 6 signal until oxygen was introduced. Thus, the 1.00 6 species is not a complex between (CH3)SSband any oxidation product such as (CH&SbO. In the quantitative run depicted in Figure 2, about 80% of the (CH3)SSbis converted to (CH:&SbO under these conditions. In support of the NMR observations, the exhaustively oxidized solution [originally mol of (CH3)3Sb]was thoroughly cencontaining 4.4X trifuged, and the clear liquid phase was carefully removed. After drying, the solid (stibinic acids) recovered amounted to only 0.0106 g, which cannot represent more than 8.71 X loT5 mol of antimony or about 20% of the antimony in the original sample. Zingaro and coworkers ( 4 , 5 )have used mercuric oxide to oxidize (CH&Sb without C-Sb bond cleavage. When we mixed red HgO and (CH&Sb in methanol in an NMR tube, no white solids were formed, and no signal a t 1.00 6 was observed during the reaction. The 0.74 6 signal disappeared and was replaced by the 1.63 6 signal during the reaction. We can account for these results by hypothesizing that oxidation by 0 2 occurs via the process outlined in Scheme 2. It is assumed that all peroxy linkages are susceptible to reduction hy excess trimethylstibine. The signal at 1.00 6 in 1130
Environmental Science & Technology
Figure 1. Proton magnetic resonance spectra (60 MHz) showing result of exposing methanol solution of trimethylstibtne to air Exposure time: A, 0 min; B, 115 min; C, 170 min; D, 300 min; E, 1400 min. Note signal on extreme left (downfield) indicated by arrow is C13-H1 satellite of methanol solvent
5.01
100
200 300 EXPOSURE
400
500
600
700
TIME ‘rninr:
Figure 2. Plot of integrated areas of PMR signals observed during oxidation of methanol solution of trimethylstibine Trimethylstibine, 0;trimethylantimony oxide, e; signal believed to be due to (CH&SbOR (R = H, OCHs), 0 . Data taken from run similar to that depicted in Figure 1. Moles of (CH3)zSbOR calculated assuming only two methyls per antimony. Estimated experimental errors represented by axes of ellipses
Figure 1 is tentatively assigned to the species (CH2)PSbOH in Scheme 2. This assignment appears likely for several reasons: If the 1.00 6 signal were due to the peroxide intermediate, it should disappear rapidly when the supply of oxygen is stopped. The persistence of the 1.00 6 signal in the absence of oxygen rules out this assignment. Compounds of the type (CH3)ZSbXreact rapidly with oxygen to give products of the type (CH3)2Sb(O)X ( 1 4 ) .The 1.00 6 signal is not observed in the oxidation of (CH&Sb with HgO where no C-Sb bond cleavage occurs. The chemical shift (1.00 6) is in the range expected for a compound of the type (CH3)PSbOH. For example, (CH3)PSbClabsorbs at 1.27 6, CH3SbC12 at 1.63 6, and all pentavalent methylantimony compounds absorb at lower field: (CH&$bXP, X = F 1.71 6, X = C12.34 6, X = Br 2.62 6, X = I 2.98 6; (CH3)dSbI 1.63 6 (3, 15, 16). A similar pattern develops for methylarsenic compounds: (CH3)4As2 1.11 6; (CH:&AsXz, X = F 2.17 6, X = C1 3.00 6, X = Br 3.23 6; (CH3)zAsOzH 1.90 6 (17, 18).
I
-T.
!
Primary Oxidation
Rearrangement
Peroxide Intermediate
II
(CH,),SbOOCH,
I -D
II
"(CH,),Sb
t
=
0"
I
(CH,),SbOH
U -
0,
I 2.0
+
~
s
-
+~ 1D
Figure 3. Atmospheric oxidation of trimethylarsine in methanol solu(CH3
(Insoluble Polymers) Rearrangement Tertiary Oxidation
(Insoluble)
To supplement the NMR scale reactions, large samples of (CH3)3Sb in diethyl ether were prepared by the procedure of Doak et al. (13). Oxidation of the solutions under protic or anhydrous conditions as outlined in the Experimental section produced high yields of trimethylantimony oxide which was isolated by conversion to the dichloride. The air oxidation of solutions of trimethylarsine in methanol was followed by the same NMR technique used in the antimony system. From the NMR spectra, it was estimated that in 48 h of exposure, less than 50% of the (CH3)3As (0.91 6) in a 0.5 M methanol solution was converted to oxidation products. As shown in Figure 3, three products are observed: (CH&AsO (1.73 6), (CH&As02H (1.90 6), and a signal at 1.26 6 tentatively assigned to (CH&AsOH. The signal for (CH3)3AsO is assigned by comparison with (CH3)3AsO prepared in methanol by oxidation of (CH3)3As with mercuric oxide ( 4 , 5 ) .Cacodylic acid (CH&AsOzH from a commercial source gave a signal at 1.90 6 in methanol. As might be expected from the introductory remarks about the acid and base properties of cacodylic acid, we found considerable variation in its signal with pH. In the presence of strong acid, it was observed a t about 2.37 6, while in strong base it was observed at 1.58 6. In Figure 3, spectra D and D' show that deliberate addition of water to the reaction mixture caused the 1.90 6 signal to be replaced by a signal at 1.85 6. Water absorbed from the air had a similar effect. A scheme for oxidation of (CH3)3As similar to Scheme 2 for oxidation of (CH3)$3b seems likely. The oxidation of (CH3)sAs with HgO in methanol was observed by PMR several times over a period of hours. Only signals for (CH&AsO and (CH&As were observed. When the partially reacted mixture was thoroughly centrifuged and the clear liquid transferred to a clean NMR tube without exposure to air, mercury (9 X loT6mol) was deposited from the solution [containing originally 2.7 X l o p 4 mol of (CH&As] over a period of several hours. These results suggest that an unstable, soluble complex of trimethylarsine and HgO [analogous to the
tion Exposure time: A, unexposed; B and B', several hours; C, several days: D and D', more than a week. B' is tenfold increase in spectrum amplitude of B. D' shows result of adding drop of water to D. Signal at extreme left in spectra A-D is C13 satellite of CH,OH solvent. More than half of original (CH&As in sample estimated to be lost by evaporation during oxidation process. (Note relative intensities of CHBOHC13 satellite in spectra A and D)
well-known mercury halide complexes (19)] is formed in the presence of excess (CH3)3As. The oxidation reactions discussed above do not easily lend themselves to quantitative study. Nonetheless, we can deduce important information about the rates of the reactions without being concerned about their detailed stoichiometry, if we are willing to accept calculations which can presently permit only order of magnitude estimates of the rate constants. A mathematical procedure which gives reasonable results is described in the following approach. The oxidation of arsines and stibines in the liquid phase by atmospheric oxygen is a heterogeneous reaction system which can be represented in simplified form by equations which assume a second-order kinetic rate-controlling step. ki
-
Oz(atmospheric) T-, Oz(disso1ved) k-1
(CH3)aE
+ Oz(disso1ved)
kz
fast
[(CH3)3E02]-+products
In such a system, the reaction rate can be limited either by the rate of diffusion of oxygen into the liquid phase (reaction in the diffusion controlled regime) or the rate of combination of the reactants in the homogeneous liquid phase (reaction in the kinetic-controlled regime) (20). In addition, there is a transition regime where the rate.of diffusion approaches the rate of combination. For these systems, we may state three general conditions: Diffusion controlled Transition Kinetic controlled
h2[02(d)][(CHs)aE]>> k 1 [ 0 2(atm)] h d O ~ ( d ) l [ ( C H d ~=Ekl 1 [ 0 2 (atm)] 500 min), the reaction has moved out of the diffusion-controlled regime and into the kinetic-controlled regime. However, it is possible that the secondary and tertiary oxidation processes (Scheme 2) are fast enough relative to the primary oxidation process to effectively compete with (CH3)3Sbfor infusing oxygen and thus attenuate the rate at which (CH&Sb is consumed. The observation that a secondary oxygen scavenger, which we believe to be (CH:J:$bOH absorbing a t 1.00 6, does accumulate in the liquid phase during the early part of the reaction and is not consumed until the concentration of (CH3):ISbreaches a low value (exposure time > 700 min) would seem to weigh against this intermediate being a major oxygen sink via secondary oxidation. Although we have not demonstrated experimentally the accumulation of oxygen in the liquid phase in the latter part of the reaction (exposure time > 500 min), we can make the joint arguments that transition conditions do not exist before [(CH:&jSb]is reduced to about 10-l M, and a t that time the oxygen concentration must be less than the saturation concentration 10-2 M. In this way we have chosen maximum values for both [(CH3)3Sb]and [02(d)]under transition conditions and we can calculate that k2 for reaction of oxygen with M-'.s-': (CH:I):$b in methanol is greater than
-
h 2 > 10-5 M.s-'/(lO-l M.10-2 M) = 10-3 M-I.s-1 1132
Environmental Science & Technology
Before leaving this subject, we might note that the formation of solid products of secondary and tertiary oxidation only seems to become very important during the latter stages of the experiment when the reduction of the hypothesized peroxylinkages is suppressed for lack of a suitable reducing reagent in high concentration. In the early part of the experiment, the excess (CH&Sb played the part of the reducing reagent (Scheme 2). The oxidation of (CH&As is much slower than oxidation of (CHz3)$3b.As a matter of fact, the (CH&As tended to evaporate when exposed to air faster than oxidation products accumulated in the liquid phase. Clearly the oxidation reaction was not limited by diffusion of oxygen into the liquid phase. Although the observations are rather crude, it is estimated from the results in Figure 3 that for (CH&As oxidation in methanol k.2 < M-l s-l. The slowness of the oxidation of (CH:3)3Asis contrary to the impression given in many descriptions of alkylarsines. However, this result is in keeping with the fact that Challenger ( 2 )has shown that (CH&As can be isolated from highly aerated biological systems. The reputation for rapid oxidation of alkylarsines no doubt arises from work with the higher analogues. In the case of antimony, trimethylstibine reacts more slowly than triethylstibine, and this order probably exists in the arsenic analogues (9).The usual citation that higher alkyl analogues of stibine and arsine are less reactive than the lower analogues is probably true, if the ethyl analogue is taken as the archetypal compound. In direct competition experiments in which mixtures of (CHz)$3b and (CH3)3As in methanol were exposed to air, trimethylstibine reacted 6-8 times as fast as trimethylarsine (Figure 4). This result, when coupled with the estimates of the independent absolute rates, suggests that the rate constant for (CHs)3Sb is on the order of 10-1-10-2 M-l s-l, whereas the rate constant for (CH3)sAs is on the order of 10-2-10-3
-_-
I
'I
-
I
~~
20
s
I
1.0
Figure 4. Competitive oxidation of trimethylstibine and trimethylarsine Spectrum A, unexposed samples with As:Sb, 1.0:0.8. Spectrum Band B', after 50-min exposure shows even though initial amount of (CH&Sb less than initial amount of (CH&As, total (CH&Sb oxidation products at least three times more abundant than total (CH&As oxidation products. Spectrum C, after 250 min shows more than half (CH&Sb oxidized, and only a few percent of (CH&As oxidized. Signals in Spectrum C labeled according to our tentative identification of products which, except for (CH3)3Sb0and ( C H 3 ) 3 A ~(R 0 = CH3), are not unequivocally established
M-l s-l. However, there is the possibility of piracy of the reactive intermediate hypothesized for the trimethylstibine oxidation by trimethylarsine which could decrease the difference between the rates in such a competitive oxidation (22, 231, [R3Sb02]
+R~As
+
RsSbO
+ R~AsO
In view of the fact that an important amount of both (CH3)3Sb and (CH3):jAscan volatilize into ambient air from liquid phases as the potent odors of their solutions demonstrate, we attempted to examine the oxidation of these compounds in the gas phase. Trimethylstibine, as noted by many authors ( I O ) , ignited over a large range of partial pressures when mixed with oxygen. At rather low pressures, as reported by Bamford and Newitt ( 9 ) ,the reaction is rapid, producing a white smoke, but ignition does not occur. An analysis similar to that described for the liquid-phase systems can be applied to Bamford and Newitt’s data. They controlled the rate of oxygen inlet (0at 7 X torr/s. (Note that these workers measured pressure in mm of sulfuric acid which we have converted to torr.) Transition regime conditions existed when the oxygen pressure was on the order of lo-* torr, and the (CH&Sb pressure was about 7 X torr (see Bamford and Newitt’s data, plot 2, Figure 5 ) . Thus, if we can assume that the reaction of (CH&Sb with 0 2 in the gas phase is primarily homogeneous and second order, k2
= (7 X
torr/s)/[(10-2torr)(7 X torr)] = 10-1 t0rr-ls-I
and since 1 torr = 5.5 X 10-5 mol/l. at 298 K, h2 = 1 X 10-*/5.5 X l o p 5 M-l s-l = 2 X 103 M-l s-l
In contrast, the apparent second-order rate constant for oxidation of trimethylarsine is on the order of M-l s-l from the data in the experimental section. Note that even if the laboratory gas-phase oxidations cited above are complicated by wall effects (chain initiation, termination, etc.) (24),the rate constants calculated above can be used to set an upper limit on the homogeneous, bimolecular reactions. Environmental systems (e.g., urban atmospheres, with the possible exception of atmosphere with high aerosol content) typically have large volume-to-surface-area ratios, and the homogeneous reaction would appear likely to dominate (25). Thus, these results would appear applicable to environmental situations. Conclusions Trimethylarsine and trimethylstibine can be regarded as environmentally mobile repositories of arsenic and antimony. The current oxidation studies provide insight into the potential for environmental dispersion of arsenic and antimony via these compounds. The observation that trimethylarsine reacts slowly with oxygen under normal conditions of concentration and temperature suggests that this compound can travel considerable distances without undergoing chemical change in aerobic systems. This result coupled with the high volatility of trimethylarsine (vapor pressure 322 torr at 298 K ) (26),compared to the expected low volatility of its oxidation products, explains how harmful and even lethal concentrations accumulated in the air of poorly ventilated rooms when wall plasters and pigments containing inorganic arsenic came under conditions suitable for microbial growth ( 2 ) .This situation was responsible for well-documented, fatal arsenic poisonings as recently as 1930 (2, 2 7 ) . The use of arseniccontaining building material has been largely discontinued in recent years. Biological methylation of antimony has not been demonstrated. However, there is no obvious thermodynamic or ki-
netic barrier to biomethylation, and the chemical similarities between antimony and Sn, Pb, As, Se, and Te, which literally surround antimony in the periodic table and all of which have been shown to be subject to biomethylation ( I , 2, 27-31), would suggest biomethylation pathways for antimony. This possibility assumes an important magnitude when the current extensive use (-lo6 kg/yr) of inorganic and organic antimony compounds in conjunction with halocarbons in fire retardant systems is recognized (32-34). Products containing antimony-based fire retardant systems include plastics, textiles, elastomers, paper, wood, paints, and coatings (33). Our current results indicate that even if antimony is subject to biological methylation, the lower volatility of trimethylstibine (vapor pressure 103 torr a t 298 K) (26) and its more rapid oxidation as compared to trimethylarsine make it very unlikely that health hazards similar to those associated with arsenic-based plasters and pigments used before 1930 could develop in aerobic environments. Nonetheless, if methylation of antimony should occur during the biodegradation of discarded or poorly maintained consumer items protected with antimony compounds, the antimony could be put into a much more water-soluble form [(CH3)3SbO] which would allow leaching of the antimony, reducing the flame retardancy of the material and contributing to pollution of waterways. Fortunately, many of the current applications of antimony in flame retardant systems for synthetic materials have the antimony dispersed throughout the material and thus effectively incapsulated so that it is only exposed during the combustion process. Other materials such as natural fibers are usually protected by “adding on” a finishing formulation containing antimony. In these systems the antimony is more exposed to biological, chemical, and mechanical separation from the material. For example, the fire retardant features of treated canvas are significantly deteriorated by exposure, implying loss of the fire retarding agent (32).At the moment, the obvious advantages of the antimony fire retardant systems seem to heavily outweigh any disadvantages, but manufacturers should be encouraged to recognize antimony as a potential hazard. In particular, the antimony oxide used in flame retarding systems should be very free of arsenic oxide which is a potential contaminant in commercial grades.
Literature Cited (1) Wood, J . M., Science, 183,1049-52 (1974).
(2) Challenger, F., Chem. Reu., 36,315-61 (1945). (3) Parris, G. E., Brinckman, F. E., J . Org. Chem., 40, 3801-03 (1975). (4) Chremos, G. N., Zingaro, R. A., J . Organometal. Chem., 22, 637-45 (1970). (5) Morris, W., Zingaro, R. A., Laane, J., ibid.,91, 295-306 (1975). (6) Merijanian, A., Zingaro, R. A., Inorg. Chem., 5,187-91 (1966). ( 7 ) Goddard, A. E., Yarsky, V. E., J . Chem. Soc., 1928, p p 719-23. (8) Dyke, W.J.C., Jones, W. J., ibid., 1930, p p 1921-27. (9) Bamford, C. H., Newitt, D. M., ibid., 1946, pp 695-701. (10) Doak, G. O., Freedman, L. D., “Organometallic Compounds of Arsenic, Antimony, and Bismuth”, Wiley, New York, N.Y., 1970. (11) Nylen, P., 2. Anorg. Allg. Chem., 246,227-42 (1941). (12) Trotter, J., Zobel, T., J . Chem. Soc., 1965, pp 4466-71. (13) Doak, G. O., Long, G. G., Key, M. E., Inorg. Syn., 9, 92-97 (1967). (14) Morgan, G. T., Davies, G. R., Proc. Roy. Soc. A , 110, 523-34 (1926). (15) Weingarten, H., van Wazer, J. R., J. Am. Chem. Sac., 88,2700-02 (1966). (16) Long, G. G., Moreland, C. G., Doak, G. O., Miller, M., Inorg. Chem., 5,1358-61 (1966). (17) Harris, R. K., Hagter, R. G., Can. J . C h e m . , 42, 2282-91 (1964). (18) Moreland, C. G., O’Brien, M. H., Douthit, C. E., Long, G. G., Inorg. Chem., 7,834-36 (1968). (19) Cotton, F. A., Wilkinson, G.. “Advanced Inorganic Chemistrv”. 3rd ed., p 520, Wiley, New York, N.Y., 1972. (20) Frank-Kamenetskii, D. A,, “Diffusion and Heat Transfer in Chemical Kinetics”, 2nd ed., pp 1-156, J. P. Appleton, Transl. Ed., Plenum, New York, N.Y., 1969. I
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(21) Clever, H. L., Battino, R., “The Solubility of Gases in Liquids”, in “Solutions and Solubilities, Part I”, M.R.J. Dack, Ed., pp 379-441, Wiley-Interscience, New York, N.Y., 1975. (22) Foote, C. S., Peters, J. W., J. Am. Chem. Soc., 93, 3795-96
(1971). (23) Rauhut, M. M., Currier, H. A,, J . Org. Chem., 26, 4626-28 (1961). (24) Melville,H., Gowenlock, B. G., “Experimental Methods in Gas Reactions”, pp 358-63, MacMillan, London, England, 1964. (25) Tuesday, C. S., Ed., “Chemical Reactions in Urban Atmospheres”, American Elsevier, New York, N.Y., 1971. (26) Rosenbaum, E. J., Sandberg, C. R., J . Am. Chem. SOC.,62, 1622-23 (1940). (27) Challenger, F., Chem. Ind. (London), 54,657-62 (1935). (28) Jernelov, A., Martin, A., Ann. Reu. Microbiol., 61-77 (1975). (29) Wong, P.T.S., Chau, Y. K., Luxon, P. L., Nature, 253,263-64 (1975). (30) Jarvie, A.U’.P.,Markell, R. N., Potter, H. R., ibid., 255,217-18 (1975).
(31) Huey, C., Brinckman, F. E., Grim, S., Iverson, W. P., Proc. Internat. Conf. on Transport of Persistent Chemicals in Aquatic Systems, Nat. Council of Canada, pp 11-73-11-78, November
1975.
(32) Lyons, J. W., “The Chemistry and Uses of Fire Retardants”, pp 209-18, 312-15, 330-32, Wiley-Interscience, New York, N.Y., 1970. (33) Little, R. W., “FlameproofingTextile Fabrics, ACS Monograph 104”, pp 358-69, Reinhold, New York, N.Y., 1947. (34) ThiBry, P., “Fireproofing”, Elsevier, Amsterdam, The Nether-
lands, 1970.
Received for reuieu: March 8, 1976. Accepted April 30, 1976. Work supported in part by the Office of Air and Water Measurement, N R S , and G.E.P. by a NRC-NBS Postdoctoral Research Associateship, 1974-5.
Laser Absorption Spectroscopy: Method for Monitoring Complex Trace Gas Mixtures Byron D. Green and Jeffrey 1. Steinfeld’ Department of Chemistry, Massachusetts Institute of Technology, Cambridge, Mass. 02139
A frequency stabilized COz laser was used for accurate determinations of the absorption coefficients of various gases in the 9-11-pm region. The gases investigated were representative of the types of contaminants expected to build up in recycled atmospheres. These absorption coefficients were then used in determining the presence and amount of the gases in prepared mixtures. The effect of interferences on the minimum detectable concentration of the gases was measured. The accuracies of various methods of solution were also evaluated.
Workers in closed environments such as submarines and spacecraft are subject to recycled atmospheres which may contain a complex mixture of species in low concentrations. The presence of other toxic substances may lower the level at which any contaminant becomes a health hazard, or the combination itself may be more dangerous than any one individual component ( I , 2 ) . As a result, atmospheric monitoring systems must be able to detect not just one or two contaminants, but many trace species all present at once in greatly different concentrations. A similar problem arises in the detection of hazardous gases in industrial environments (3). Monitoring of the attenuation of radiation over long absorption paths is the most sensitive method of detection (4-9). It lends itself well to shorter absorption paths, especially in the infrared where molecular absorptions are strong. Most polyatomic molecules have absorption bands in the near or middle infrared corresponding to a change in vibration and rotation of the molecule. The frequency pattern characteristic of each compound is unique and is routinely used to identify the molecule. Traditional broadband infrared spectrophotometric sources require frequency dispersing devices which generally have poor resolving characteristics, resulting in a blurring of the molecular “fingerprint”. A laser, in contrast, is monochromatic, easily collimated, and intense, that is, it is spectrally brilliant. A carbon dioxide laser, in addition, has the ability to emit radiation at many different frequencies, corresponding to the P and R branches of the (00’1-10”0) and (00°1-0200) molecular transitions. 1134
Environmental Science & Technology
In this report we describe the use of a COz laser system for monitoring gas mixtures under typical environmental conditions. Accurate absorption coefficients were determined for a number of species, and complete mixtures were prepared and analyzed. Detection limits in the ppm range could be obtained for many of the gases studied, for a IO-m absorption length.
Experimental The 52 strongest lines of the COSlaser, each having an intensity greater than 200 mW, are used for these experiments. The gain medium was a water-cooled, low-pressure flowing gas electrical discharge. A current regulated IO-kV power supply provides the electrical excitation. The laser cavity is composed of a curved mirror (f = 5 m) and a grating blazed for 10 p (Bausch & Lomb). T o ensure that the COZ laser frequency is centered on the molecular transition and not shifted by the cavity modes, the mirror is mounted on a piezoelectric transducer (Burleigh PZ-90). The sine wave signal generated by a lock-in amplifier (PAR Model 120) is amplified using a KEPCO OPS 2000 Op-Amp and applied to the transducer. The mechanical oscillations of the mirror produce fluctuations in the laser intensity as detected by a InSb detector (Optoelectronics Mullard type ORP-10) looking a t the laser output off the normal of the grating. The AC signal produced by the detector is proportional to the slope of the gain curve and is smallest when the mirror’s oscillation (which determines the frequency of the cavity mode) is centered on the maximum of the gain curve (which is the center of the molecular transition frequency). The magnitude of this AC signal is the output of the lock-in amplifier when the detector signal is the input. The output of the lock-in is introduced into the high-voltage KEPCO Op-Amp to produce a bias voltage on the transducer sufficient to center the mirror on the maximum of the gain curve. The sine wave voltage applied to the transducer results in -4 MHz (1.3 X lop4ern-') fluctuations of the laser frequency around the molecular transition frequency. The grating permits selection of a single lasing frequency. No further frequency dispersing or bandwidth compensating elements are required in the detection setup. An intracavity iris near the mirror is used to adjust laser power and confine lasing to the lowest order radial mode, which has a single gaussian
