Environ. Sci. Technol. 1907, 21, 102-105
Determination of Hydroxymethanesulfonate in Wet Deposition Samples Carolina C. Ang,$ Frank Llparl,"9zand Stephen J. Swarlns Environmental Science Department and Analytical Chemistry Department, General Motors Research Laboratories, Warren, Michigan 48090
A new method for determining the hydroxymethanesulfonate ion (HMSA) in wet deposition was developed by measuring the total and free formaldehyde content of the sample. Total formaldehyde (HCHO) is measured by dissociating the HMSA into formaldehyde and bisulfite with a strong base and reacting the formaldehyde with 2,4-dinitrophenylhydrazine (DNPH). The formaldehyde-DNPH derivative formed is then determined by high-performance liquid chromatography. Free formaldehyde is similarly determined in another aliquot, but without the addition of the base. The difference between the total and free HCHO measurements gives the bound HCHO or HMSA content of the sample. The method was linear for HMSA concentrations ranging from 0.06 to 6.3 pg/mL. Recovery of HCHO from the dissociated HMSA ranged from 90 to 95%. The detection limit of the method on the basis of a signal to blank ratio of 3 was 0.01 pg/mL. The method was used to measure HMSA in rain and snow samples.
Introduction Formaldehyde and other carbonyl compounds are known to form adducts with sulfite and bisulfite ions in solution. The presence of formaldehyde and sulfur dioxide in the atmosphere coupled with their high solubility in water suggests that hydroxymethanesulfonate(HMSA) can form in atmospheric droplets at low pH (I,2). The formation of HMSA can be generally expressed as HCHO + S(IV) F? HCHO-S(1V) adduct. Dasgupta et al. (3),in their study of the equilibrium constant for the dissociation of HMSA, found that the adduct is most stable between pH 4 and pH 5, which is within the typical pH range of acidic wet deposition. In their study of cloud and fog chemistry in Los Angeles, Richards et al. (4) measured S(1V) concentrations in cloud water and found that the concentrations were more than 100 times greater than those calculated from the ambient SOz concentration, the solution pH, and the acid dissociation constant for sulfurous acid. They estimated that approximately one-third of the S(IV) in the condensed phase was in the HMSA adduct form; this was based on the measurement of the free HCHO content of their samples. Similarly, Klippel and Warneck (5) found that the concentration of formaldehyde in aerosol was about 1000 times higher than the predicted equilibrium concentration between the gas-phase formaldehyde and the aqueousphase formaldehyde of the aerosol. They attributed this enrichment to the formation of the formaldehyde addition compounds. The determination of HMSA is needed to assess the importance of HMSA to droplet acidity and to describe the atmospheric sulfur budget. HMSA has been determined by ion chromatography (3,6,7).However, probbems associated with poor resolution and dissociation of the adduct were encountered. Recently, Munger et al. (2) reported the determination of HMSA in fogwater samples by a novel ion-pairing chromatographic technique.
* Environmental Science Department, 8 Analytical
102
Chemistry Department.
Environ. Sci. Technol., Vol. 21, No. 1, 1987
This paper describes a simple procedure for the indirect determination of HMSA by measuring the total and free formaldehyde contents in the sample. The method assumes that all the bound HCHO is in the HMSA adduct form. Although we recognize that HCHO-amine adducts can also form in solution (a), there are no reported measurements of such species in wet deposition samples. The method was used to analyze HMSA in wet deposition samples obtained from a site in suburban Detroit.
Experimental Section Standards and Reagents. Sodium hydroxymethanesulfonate of 95% purity was purchased from Aldrich Chemical Co. (Milwaukee,WI) and from Chem Service Co. (West Chester, PA). Standard solutions of HMSA were prepared by dissolving the sodium salt of the compound in distilled water. The formaldehyde standard solutions were prepared by diluting a previously assayed 37% formalin solution (J. T. Baker, Phillipsburg, NJ) with acetonitrile (9). The 2,4-dinitrophenylhydrazine(DNPH) was obtained from J. T. Baker and recrystallized twice from high-performance liquid chromatography (HPLC-) grade acetonitrile to remove carbonyl impurities. Standard solutions of DNPH were then prepared in acetonitrile. All other reagents were prepared by dissolving the required amount of J. T. Baker reagent-grade chemicals in HPLC-grade water. The mobile phase was generated with HPLC-grade solvents obtained from Fisher Scientific (Fair Lawn, NJ). Instrumentation. A microprocessor-controlled Varian 5000 liquid chromatograph (Palo Alto, CA) equipped with a Perkin-Elmer LC-85 spectrophotometric detector (Norwalk, CT) set at 365 nm was used for the analysis of HCHO. A Valco six-port air-actuated valve (Valco Instrument, Houston TX) was used to inject 30 pL of sample. A 4.6 mm i.d. X 25 cm Zorbax ODS column was used for the analysis. The mobile phase utilized was 55/45 acetonitrile/water flowing at 1.0 mL/min. Total S(1V) and SO-: were analyzed with a Dionex Model 10 ion chromatograph (Sunnyvale, CA) equipped with a conductivity detector and an AS4 anion column (Dionex HPIC) followed by a fiber suppressor. The eluent used was 2.8 mM HC03-/2.3 mM CO:-. The sample pH was measured with an Orion combination electrode and pH meter. Analytical Procedure. HMSA is indirectly determined by measuring the total and the free formaldehyde concentrations in the sample. The total formaldehyde, which consists of both free and adduct-bound formaldehyde, is measured after dissociating the HMSA into formaldehyde and bisulfite with the addition of a strong base at pH 12-13. The formaldehyde in the solution is then reacted with DNPH to form the HCHO-DNPH derivative, which is determined by high-performance liquid chromatography (HPLC). Free formaldehyde is similarly determined in another aliquot, but without the addition of the base. The difference between the total and the free formaldehyde measurements gives the bound formaldehyde or HMSA content of the sample. (A) Total Formaldehyde. One milliliter of the sample was pipeted into a small graduated tube, and 50 p L of 1 N NaOH was added from a dispensing pipet to dissociate
0013-936X/87/0921-0102$01.50/0
0 1986 American Chemical Soclety
100
B2
2 Z
95
n
n I
90
0-
I 0 I
I V
I
9 w_
9
85
be
80
w *
> K
75
5
10
15
20
MOLAR RATIO DNPH : HS03Figure 1. Effect of HS0,-/HCHO mation.
molar ratio on HCHO-DNPH for-
the HMSA into HCHO and bisulfite. The tube was capped and inverted twice to mix. Next, 100 pL of 1N HC104 was added to acidify the solution and to catalyze the HCHO-DNPH reaction. After mixing by inverting the tube, the sample in the tube was then made up to 2 mL with 1.6 X M (0.31 mg/mL) DNPH reagent. The tube was capped and inverted several times. The solution was allowed to stand for at least 45 min to complete the DNPH derivatization reaction. A 30-yL aliquot was then injected into the HPLC for analysis. (B) Free Formaldehyde. The procedure was the same as that for total formaldehyde determination except that no NaOH was added to dissociate the HMSA compound. Wet Deposition Sample Collection. Sampling of wet depositions ran from September 1983 to March 1984. The rain and snow samples were collected in polyethylene buckets with an Aerochem-Metrics wet/dry sampler (Miami, FL). The sampling site is located on a 330-acre vacant parcel of land in Warren, a suburb of Detroit, MI. The samples were filtered through a 0.45-pm Nucleopore filter the same day and refrigerated without further treatment, until analysis. Measurement of the samples for HCHO, and consequently HMSA, was completed within 5 days of collection. The total S(1V) concentration in four of the rain samples was determined within a day after collection.
Results and Discussion HMSA Dissociation and DNPH Reaction. The determination of free HCHO relies on the assumption that bound formaldehyde in the HMSA compound does not react with DNPH to form the hydrazone derivative. This was verified by experiments in which standard HMSA solutions were added to the DNPH reagent at pH 2, and no HCHO-DNPH derivatives were formed. The addition of a strong base to the HMSA solution dissociates the adduct into free formaldehyde and bisulfite. We found that complete dissociation of the adduct is achieved by adding 1 N NaOH to pH 13. The reaction of formaldehyde with DNPH occurs in an acidic medium (10). Therefore, acidification of the solution was necessary. This was accomplished by the addition of 1 N perchloric acid until a pH of 1-2 was obtained. We found that the acidification process did not cause re-formation of the adduct. A t pH 2, the reported dissociation constant for HMSA is 1.54 x loy4,which is an order of magnitude greater than that at pH 4 and 5 (3). The reaction rate of HCHO and S(1V) in solution is also slower at this lower pH (11).
Figure 2. Optimization of DNPH reagent in the presence of HS03-.
Table I. Recovery of Total HCHO in Spiked Rain Samples sample"
A B C D
E F
total HCHO, fig/mL calcd DNPH method 0.154 0.183 0.120 0.120 0.651 0.396
% recovery 98 95.6 101.7 105 96.8 98.5
0.151 0.175 0.122 0.126 0.630 0.390
mean: 99.3 SD: 3.5 "Samples A-D were spiked with 0.10 pg/mL HCHO. Samples E and F were spiked with 0.38 pg/mL HCHO.
In order to evaluate the effect of the molar ratio of bisulfite to formaldehyde on HMSA dissociation, we added variable amounts of bisulfite to a constant amount of HCHO (0.13 pmol) and determined the total HCHO concentration using the previously described procedure (Figure 1). We found that at a HSO,-/HCHO molar ratio of up to 2 over 95% of the liberated HCHO reacted with the DNPH reagent to form the HCHO-DNPH derivative. When the molar ratio was greater than 2, partial re-formation of the HMSA occurred and interfered with the analysis. This effect was also seen by Du Val et al. (12). However, this was not a problem in the analysis of our wet deposition samples where the HSO,-/HCHO molar ratio was less than 2. Reagent and Reaction Time Optimization. The amount of DNPH reagent required to produce a quantitative yield of the HCHO-DNPH derivative was determined by adding varying amounts of DNPH to a fixed amount of HCHO and HSO; in solution at pH 2. Figure 2 shows that over 95% of the total HCHO in solution is derivatized at a DNPH/HSO< molar ratio of 10. In another series of experiments, the optimum reaction time for the HCHO-DNPH derivatization reaction was determined to be about 40-45 min (Figure 3). Calibration and Method Validation. Standard solutions of HCHO were used to prepare a calibration curve. The detection limit of the method on the basis of a signal to blank ratio of 3 was 0.01 yg of HCHO/mL. The results from the determinations of known concentrations of HMSA solutions were over 90% of their theoretical values. As a linearity check, standard HMSA solutions ranging in concentration from 0.06 to 6.36 yg/mL were measured and found to be linear in this range. Six rain samples spiked with HCHO were also analyzed and showed a better than Environ. Sci. Technol., Vol. 21, No. 1, 1987
103
Table 11. Concentrations of Various Species in Rain and Snow total HCHO, wLg/mL
free HCHO, wg/mL
bound HCHO, wg/mL
ND
ND
ND
ND
0.31 0.05 0.05 0.20 0.20 0.19
0.23
0.08 0.05
HMSA, rg/mL"
so:-, PH
wg/mL
4.1 4.3 4.3 4.7 5.5 5.3 3.9 3.7 4.0 4.2 4.2 4.1 4.7 4.5 4.1 6.1
3.9 8.2 1.6 1.5 9.6 2.6 5.1 1.5 5.2 2.2 3.7 4.2 2.7 1.7 3.3 1.7
4.7 4.2 4.2 4.6 3.9
1.5 1.7 1.5 2.5 1.5
rain 9-23-83 10-02-83 10-12-83 10-13-83 10-18-83 10-24-83 11-03-83 11-07-83 11-14-83 11-16-83 11-21-83 11-22-83 11-28-83 3-12-84 3-13-84 3-15-84
ND ND
0.05
0.09 0.10 0.17
0.11 0.10 0.02
0.30 0.19 0.19 0.41 0.38 0.07
ND
ND
ND
ND
0.07 0.07 0.16 0.10 0.11 0.13 0.13 0.08
0.04 0.04 0.08
0.03 0.03 0.08 0.10 0.01 0.06 0.09 0.01
0.11 0.11 0.30 0.37 0.04 0.22 0.34 0.04
ND 0.10 0.07 0.04 0.07
snow 1-30-84 2-06-84 2-13-84 2-27-84 2-28-84
0.06
0.01
0.05
0.19
ND ND
ND ND ND ND
ND ND
ND ND
0.03
0.11
ND
ND
0.03
ND
HMSA = (bound HCHO/M. of HCHO)
-+- --e-
*H
80-
f
/
X
r'
M , of HMSA.
-4
-0-
I I I 40- I I 60-
20 0
I
I
20
40
,
60
4
I
80
100
120
REACTION TIME (min) Figure 3. Time study for HCHO-DNPH reaction.
95% HCHO recovery (Table I). Method Application. Table I1 shows the pH, [SO?-], and the HCHO and HMSA concentrations in rain- and snowwater determined by this method. The HMSA concentrations were calculated from the bound HCHO concentrations. HMSA was found in all but two of the rain samples analyzed. The concentration ranged from a low of 0.04 pg/mL to a high of 0.41 pg/mL. Of the five snow samples analyzed, only two showed measurable HMSA concentrations of 0.11 and 0.19 pg/mL. For comparison purposes, the molar equilibrium constant ( K )for the reaction HMSA F? HCHO + HSOs was calculated for four rain samples in which the total S(1V) concentration was determined by ion chromatography. We assumed that all bound S(1V) is in the HMSA adduct form. The HS03- concentration is then obtained by difference between the total S(1V) and the bound S(1V) concentration. K for the dissociation of HMSA may be expressed as K= [free HCHO]([total S(IV)]-[bound S(IV)])/ [HMSA] The calculated K for four samples with a pH of 4-4.5 varies from 1 X lo4 to 3 X lo4 M. These K values indicate 104
Environ. Sci. Technol., Vol. 21, No. 1, 1987
adduct stability at pH 4-5 and are in reasonable agreement to those of Kok and Dasgupta (13, 3), who reported K values of lo-' and M for a pH range of 4-5, respectively. As seen in Table 11, measurable amounts of HMSA were found in the samples analyzed. Although the [HMSA] is influenced by the sample pH (2),we found no correlation between the amount of rain HMSA and the rain pH. The total sulfate concentration for these samples was determined by ion chromatography and ranged from 1.5 to 9.6 pg/mL. Comparing these sulfate concentrations to the HMSA concentrations, we calculated an average HMSA/SO?- molar ratio of 0.06. This implied that HMSA represented only a small fraction of the total sulfate concentration, and that HMSA was probably a minor contributor to the overall acidity in these samples. Kumar (14),in his modeling studies of rain composition, calculated the HMSA formation time constant required for the reaction of bisulfite and HCHO to approach equilibrium. His study showed that this reaction in the falling raindrops cannot approach equilibrium and would continue further in the collected rain sample until equilibrium is achieved. However, he also stated that, for cloud droplets whose average life-time is approximately 30 min, there is considerably more time for the reaction to approach equilibrium. Most of our rain samples have pH values of 4-5, where the HMSA is most stable and where formation rather than decomposition is favored. Therefore, it is possible that the amount of HMSA found in our rain samples represents an upper limit and may not be indicative of the actual concentration present in the raindrops.
Acknowledgments We acknowledge Peter J. Groblicki for valuable suggestions during the course of this work and William Scruggs for assistance in the experimental work. Registry NO.HMSA, 75-92-3; HCHO, 50-00-0; HzO, 7732-18-5.
Literature Cited (1) Munger, J. W.; Jacob, D. J.; Hoffmann, M. R. J . Atmos. Chem. 1984, I, 335.
Environ. Sci. Technol. 1987, 27, 105-110
(9) Reagent Chemicals; American Chemical Society: Washington, DC, 1961; Vol. I, p 222. (10) Lipari, F.; Swarin, S. J . Chromatogr. 1982, 247, 297. (11) Boyce, S. D.; Hoffmann, M. R. J . Phys. Chem. 1984,88, 4740-4746. (12) Du Val, D. L.; Rogers, M.; Fritz, J. S. Anal. Chem. 1985, 57, 1583-1586. (13) Kok, G. L.; Gitlin, S. N.; Lazrus, A. L. J. Geophys. Res. 1986, 91, 2801-2804. (14) Kumar, S. Atmos. Environ. 1986, 20, 1015-1024.
Munger, J. W.; Tiller, C.; Hoffmann, M. R. Science (Washifigton,D.C.) 1985,231, 247-249. Dasgupta, P. K.; DeCesare, K.; Ullrey, J. C. Anal. Chem. 1980,52, 1912-1922. Richards, L. W.; Anderson, J. L.; Blumenthal, D. L.; McDonald, G. L.; Kok, G. R.; Lazrus, L. A. Atmos. Environ. 1983, 17, 911-914. Klippel, W.; Warneck, P. Atmos. Environ. 1980,14,809-818. Dasgupta, P. K. Atmos. Environ. 1982, 16, 1265-1268. Lindgren, M.; Cedergren,A.; Lindgren, J. Anal. Chim. Acta 1982, 141, 279-286. Walker, J. F. Formaldehyde, 3rd ed.; Krieger: New York, 1975.
Received for review January 14, 1986. Revised manuscript received August 19, 1986. Accepted August 26, 1986.
Motor Exhaust Emissions as a Primary Source for Dicarboxylic Acids in Los Angeles Ambient Airt Klmltaka Kawamuras and Isaac R. Kaplan*
Institute of Geophysics and Planetary Physics, University of California at Los Angeles, Los Angeles, California 90024
Low molecular weight dicarboxylic acids (C2-Cl0)were analyzed in Los Angeles air and auto exhaust, as well as greenhouse air, soil, dust, and bog sediment samples, as dibutyl esters by gas chromatography and gas chromatography-mass spectrometry with fused silica capillary columns. In the Los Angeles ambient atmosphere, 19 dicarboxylic acids in the range C2-Clo were identified, including straight-chain, branched-chain, cis- and transunsaturated, and aromatic acids. Oxalic acid is the dominant species, followed by succinic, malonic, maleic, glutaric, adipic, and phthalic acids. Total concentration of C2-Clo diacids detected in the ambient atmosphere ranged from 5.5 to 21.2 nmol/m3 (average 12.3 f 6.1 nmol/m3). By contrast, gasoline and diesel exhaust samples, collected under idling conditions, showed that distributions of the diacids are similar to those of air samples, but their concentrations are 28 (gasoline) and 144 (diesel) times higher than the average concentration of atmospheric diacids. These results indicate that engine exhaust is an important source of diacids in the urban atmosphere. H
Introduction Short-chain dicarboxylic acids in the atmosphere are thought to derive from photochemical reactions, principally, oxidation of cyclic olefines and other hydrocarbons (1-3). Because these compounds are polar and also nonvolatile under ambient conditions, dicarboxylic acids may be associated with the formation of aerosols and, in part, be responsible for reduction in the visibility of urban atmospheres (Le., they may act as nuclei for aerosol formation). Thus, studies of short-chain dicarboxylic acids may be helpful to a better understanding of several problems in atmospheric chemistry. However, the analytical techniques previously used were inadequate for the accurate characterization and quantification of short-chain diacids (IC,) in the atmosphere. Dicarboxylic acids have previously been quantified in southern California aerosol samples by paper chromatography and infrared spectra (4), computerized high-resolution mass spectrometer (HRMS)-thermal analysis for 'Publication No. 2714 of the Institute of Geophysics and Planetary Physics. $ Present address: Department of Chemistry, Woods Hole Oceanographic Institution, Woods Hole, MA 02543. 0013-936X/87/0921-0105$01.50/0
semiquantitative determination of C5-C7 diacids ( 1 , 5 , 6 ) , and capillary gas chromatography-mass spectrometry (GC-MS) for saturated C3-Clo aliphatic dicarboxylic acids employing methyl ester derivatives (2).However, these techniques failed to detect oxalic acid. Oxalic acid was recently measured in upper atmospheric and stratospheric particles from Colorado by ion chromatography (3), although other diacids were not detected by this method. The literature reports indicate, therefore, that no single method is available to accurately determine the distribution of C2-Clo dicarboxylic acids in the ambient atmosphere. Here, we describe a new method to measure C2-Clo dicarboxylic acids in atmospheric aerosols. The atmospheric diacids were collected on two quartz fiber filters in series, the second being impregnated with KOH. The diacids are extracted with water, derivatized to dibutyl esters, and determined with high-resolution capillary GC and GC-MS. This technique has also been used for the analysis of motor exhaust collection in one gasoline- and one diesel-operated automobile. To understand if diacids are significantly produced in other environments, soil, urban dust, peat bog sediment, and the air from a greenhouse were also studied. Experimental Section Reagents and Materials. Pure water was prepared by oxidizing organic impurities in distilled water with KMnOJKOH, followed by distillation in glass. KOH solutions were prepared after KOH pellets, which were prewashed with methanol, were heated at 500 "C for 3 h. Other reagents and solvents were prepared as described previously (7). Quartz fiber filters (25 cm X 21.5 cm) were purchased from Pallflex Products Corp. (Putnam, CT) and were cut into 47 mm diameter discs. The trapping efficiency of 0.3-pm dioctyl phthalate on the filter quoted by Pallflex Products Corp. is 99.97 % . The filters were heated at 500 "C for 3 h to combust organic contaminants, rinsed in 1% KOH solution (to prepare KOH-impregnated filters) or pure water (to prepare neutral filters), and then dried in an oven at 80 OC. Sample Collection. Air samples were collected on both neutral and KOH-impregnated quartz filters, which are placed in combined Nucleopore filter holders (Pleasanton, CA) in series and connected to a flowmeter or a calibrated
0 1986 American Chemical Society
Environ. Sci. Technol., Vol. 21, No. 1, 1987 105
