Environ. Sci. Technol. 1990, 2 4 , 384-385
CORRESPONDENCE Comment on “Acidification of Adirondack Lakes”
Table I. Absorbance Values for “Yellow”, “Red”,and Peak Wavelengths of Methyl Orange in Dilute HCI Standards
SIR: Asbury et al. have addressed several important details necessary for a quantitative reconstruction of acidification trends in Adirondack lakes from 1929 to 1985 ( I ) . We have examined additional critical assumptions that control the feasibility and reliability of such a quantitative reconstruction, using the 1929-1934 methyl orange (MO) alkalinity determinations. In this letter we summarize these results, concluding that the historic Adirondack MO alkalinity measurements have too many important, confounded errors and are too imprecise to be used in any quantitative trend analysis. The most important source of bias accounted for by Asbury et al. is the overestimation of acid neutralizing capacity (ANC) during MO titrations ( I ) , which usually continue to a fixed pH end point that is more acid than a modern Gran potentiometric determination would be for the same lake water sample. The crux of their argument is that the MO titration end point occurs at a mean pH of 4.25 (fO.10 = 2s, n = 60, ANC = 85 pequiv/L) ( I ) compared to an end point pH of 4.04 (kO.10, n = 24, ANC = 162 f 1pequiv/L) reported by Kramer and Tessier (2). Kramer et al. (3)presented both the 4.04 and a 4.19 (based on unpublished materials) pH unit end point for correcting MO alkalinity determinations in their study for the National Academy of Sciences. Either documented approach (1,2) is affected by the accuracy and precision of the pH electrode measurement of the end point, plus the subjective choice of what is “faintest pink” (2) at the end point. Asbury et al. (1)have estimated the precision of the MO end points by a single pH measurement following 60 MO ANC titrations of one water sample by students. The ANC determined with a MO titration also depends on the accuracy of the end point pH measurement, which we estimated without pH electrodes, but was not accounted for by Asbury et al. Past quantitative investigations of pH indicator color changes have considered the sharpness of transition at the equivalence point (561, but only for concentrated buffer solutions. Using a high precision spectrophotometer and the unacidified methyl orange sodium salt ( 3 , 4 )(0.500 g/lOOO g of solution) (2) a t a concentration of 2 drops of MO (1 drop = 0.041 g f 0.007 (Bs), n = 30) per 50 mL of sample (2), we have measured the absorbance of dilute HC1 standards, made to within f0.003 pH unit (7) (except pH 5.00, within fO.01 pH unit). With this approach we can analyze the determination of end point color on a quantitative basis. Other studies rely on pH electrode measurements, which contribute an additional imprecision of at least k0.05 pH unit, 2s (7) and unknown bias. Table I summarizes the most important results from 25 absorbance spectra scanned from 200 to 700 nm with a Nicolet-Phillips Analytical Cambridge 8740 UV/vis scanning spectrophotometer to within fO.001 absorbance unit within f0.3 nm of the stated wavelength. When the sample’s pH is equal to the “apparent” dissociation constant of the indicator, the ratio [yellow form MO]/[red form MO] becomes equal to 1, and the indicator 384
Envlron. Sci. Technol., Vol. 24, No. 3, 1990
HCl std pH 2.04 3.02 4.04 4.15 4.19 4.25 4.35 5.00
absorbance “yellow” “red” 422 nm 522 nm 0,001 0.006 0.021 0.029 0.028 0.025 0.023 0.035
0.127 0.063 0.021 0.024 0.018 0.016 0.010 0.015
peak absorbance nm
Abs
ratio of yellow/red
508 506 480 474 474 474 471 467
0.135 0.070 0.042 0.050 0.051 0.044 0.041 0.052
0.008 0.1 1.0 1.2 1.6 1.6 2.3 2.3
A,
ideally will have a color due to an equal mixture of the “acid” and “alkaline” forms (8). In order to make such a comparison, concentration and absorbance must be linearly related a t a given wavelength. We have verified that the Beer-Lambert law holds at up to 5 times the MO additions used in these experiments, as earlier workers have also observed (9). From Table I, the solution with a pH of 4.04 has the ratio of “yellow” to “red” absorbance equal to 1.0, but the “relative visibility factor” (relating the normal response to radiation of equal intensity) of the human eye is 0.004 at 420 nm and 0.710 a t 520 nm ( I O ) . This means the eye is more sensitive to red than yellow solutions of equal absorbance. Thus, the pH 4.04 solution appears with a slight reddish tint, “faintest pink” ( I ) , analogous to the MO titration end point. This is in agreement with the beginning of color change for similar solutions as reported by Kolthoff and Stenger (11). A visual comparison of the pH 4.25 and pH 4.35 solutions resulted in them being classified as the same color, a faint yellow. The pH 4.19 and pH 4.15 solutions occupy a region of transition between faint yellow and faint pink, with the actual point of demarcation being subject to individual preference, training, and confounding parameters affecting the chemistry, such as the presence of dilute amounts of organic acids. Several of these sources of variability are fixed in Asbury et al.’s estimation of the MO pH end point, which may result in statistical bias. The visible absorbance spectra observed in this work did not show bimodal distributions of peak absorbance corresponding to color. Instead, one absorbance peak shifted from more “yellow”wavelengths to more “red” wavelengths as the pH decreased (Table I). Obviously, we have no way of distinguishing quantitatively which subjective color end point (hence, end point pH) different analysts actually chose during the 1920-1930s. It is likely that such early work showed at least the disagreement (imprecision) seen between the modern studies (1-3). A lesser concern is the average MO alkalinity correction reported by Asbury et al. ( I ) (-54.6 pequiv/L at pH 4.25) and Kramer and Tessier (2) (-81 pequiv/L at pH 4.04). Table I1 shows the results of our calculations for a MO alkalinity correction (for MO end point of pH 4.04), which varies with the dissolved inorganic carbon content (DIC) of a water sample. The calculations are based on Kramer et al.’s Appendix D, eqs lb, 4b, 4d, and 7g ( 3 ) . Table I1 shows a representative range of corrections
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0 1990 American Chemical Society
Environ. Sci. Technol. 1990, 2 4 , 385-387
Table 11. Variation of Methyl Orange Alkalinity Corrections with Dissolved Inorganic Carbon at pH 6"
concn, pequiv/L [Alk~o] [Alk] corrctn
DIC, mg/L
relation to DIC population of ELS-I region 1 lakesb
0.2 0.6 1.6 3.3 13.3 49.0
min. observed DIC (region 1) 20th percentile DIC (region 1A) 50th percentile DIC (region 1A) 80th percentile DIC (region 1A) 95% confid limit DIC (region 1) max. observed DIC (region 1)
96.2 106.3 131.5 174.4 426.5 1326.5
4.1 14.4
40.0 83.6 339.8 1254.6
-91.1 -91.0 -90.5 -89.8 -85.7 -70.9
"At other pH values, [AlkMo], and [Alk] values will change, but the correction factor remains constant. *References 12 and 13.
calculated for DIC concentrations measured during the EPA's 1984 Eastern Lake Survey-Phase I(I2,13). Region 1A contains the Adirondack lakes greater than or equal to 4 ha surface area. These are a subset of lakes sampled in region 1, also including Central and Southern New England, Maine, and the Poconos/Catskills. As a result, the MO alkalinity correction in the Adirondacks region should vary from about -91 to -86 pequiv/L, depending on a sample's inorganic carbon concentration, rather than the average correction of -54.6 pequiv/L used by Asbury et al. (I) (Table 11). Although the individual [AlkMo] and [Alk]values vary as a function of pH, the MO correction does not and is dominantly controlled by the variation of [DIC]. An additional complication, which we have not noticed in the modern acidification literature, concerns the inability of MO to react with weak acids, such as organic acids (4,14). Stieglitz (4)hypothesized that the red form of MO reacted as a very weak base; "red" salts of weak acids were "completely hydrolyzed and incapable of existence, the liberated base reverting ... to the stable yellow isomer. That is why methyl orange is not a sensitive indicator for weak acids". This effect could prevent the detection by MO of weak acids that contribute to a solution's ANC as determined by modern potentiometric methods, resulting in underestimates of ANC. Alternatively, these reactions could also result in a higher MO pH than a potentiometric pH in the presence of organic acids, causing more titrant to be added to the fixed end point, resulting in a higher historic ANC value than would be determined potentiometrically. Additional work is necessary in order to understand the interaction of MO with organic acids during acid titrations of dilute solutions (ref 3, Appendix D). The dissolved organic carbon content (DOC) (used as a measure of organic acid content) could be approximated as we have used the modern data from the Eastern Lake SurveyPhase I. However, this would be a poor approximation to the actual lake DOC of the 1920-1930s. The qualitative correlation between high DOC content and increasing lake color is well-known, and unfortunately, one of the qualitative observations for acidification is that some lakes have turned clear since that time, which implies past lake DOC concentrations were different from today's observations. We thank D. Heggem, US.E.P.A., EMSL-Las Vegas, for providing support and encouragement during the preparation of this manuscript. Registry No. Water, 7732-18-5; methyl orange, 547-58-0.
Literature Cited (I) Asbury, C. E.; Vertucci, F. A,; Mattson, M. D.; Likens, G. E. Environ. Sci. Technol. 1989, 23, 362-65. (2) Kramer, J. R.; Tessier, A. Enuiron. Sci. Technol. 1982,16, 606A-15A. (3) Kramer, J. R.; Andren, A. W.; Smith, R. A.; Johnson, A. H.; Alexander, R. B.; Oehlert, G. In Acid Deposition: Long Term Trends; National Research Council; National A c a 0013-936X/90/0924-0385$02.50/0
demy Press: Washington, DC, 1986; p 231. (4) Stieglitz, J. J. Am. Chem. Soc. 1903, 25, 1112-27. (5) Cacho, J.; Nerin, C.; Ruberte, L.; Rivas, E. Anal. Chem. 1982,54, 1446-49. (6) Bhuchar, V. M.; Kukreja, V. P.; Das, S. R. Anal. Chem. 1971,43, 1847-53. (7) Metcalf, R. C. Analyst 1987, 112, 1573-77. (8) Bassett, J.; Denney, R. C.; Jeffery, G. H.; Mendham, J. Vogel's Textbook of Quantitative Inorganic Analysis including Elementary Instrumental Analysis; Longman: London, 1978. (9) Sidgwick, N. V.; Worboys, W. J.; Woodward, L. A. Proc. R. SOC.London 1930, A129, 537. (10) Born, M.; Wolf, E. Principles of Optics, 2nd ed.;Macmillan: New York, 1964; p 185. (11) Kolthoff, I. J.; Stenger, V. A. Volumetric Analysis, 2nd ed.; Wiley: New York, 1947; Vol. 11, pp 57-62. (12) Linthurst, R. A.; et al. Characteristics of Lakes in the Eastern United States. Volume 1: Population Descriptions and Physico-Chemical Relationships; EPA600/486/007a; U.S. Environmental Protection Agency: Washington, DC, 1986. (13) Overton, W. S.; et al. Characteristics of Lakes in the Eastern United States. Volume 11 Lakes Sampled and Descriptive Statistics for Physical and Chemical Variables; EPA600/4-86/007b; U.S. Environmental Protection Agency Washington, DC, 1986. (14) Windholz, M., Ed. The kferck Index, 10th ed.; Merck & Co.: Rahway, NJ, 1983; p 874.
Richard C. Metcalf," Robert W. Gerlach
Environmental Chemistry Assessment and Quality Assurance Departmentt Lockheed Engineering & Sciences Company 1050 E. Flamingo Road, Suite 209 Las Vegas, Nevada 891 19 'Although the information in this letter has been funded in part by the United States Environmental Protection Agency under Contract 68-03-3249 to Lockheed Engineering & Sciences Co., it has not been subjected to Agency review. It therefore does not necessarily reflect the views of the Agency and no official endorsement should be inferred. The mention of trade names or commercial products does not constitute endorsement or recommendation for use. This work is a contribution to the United States Environmental Protection Agency's Aquatic Effects Research Program, which is a part of the National Acid Precipitation Assessment Program.
SIR: Asbury et al. (I) recently gave an assessment that suggests that there has been a median decrease in alkalinity of -50 pequiv/L from the 1930s to the 1970-1980s for surface waters of the Adirondack region of New York. They contrast their results to those reached in a National Research Council (NRC) report (2),which concluded that there was a median loss of alkalinity between 0 and 44 pequiv/L. The Asbury et al. report also makes various claims to suggest that the "no change" conclusion of the NAS report is wrong. I disagree with this conclusion for the reasons stated below. I also restate important points of the NRC report that were not considered in the Asbury et al. report. First, the problem of making a clear statement on change in alkalinity (actually acid neutralizing capacity) or change in pH for data in the 1930s is difficult because the unknowns in the colorimetric analyses are as large or larger than inferred change. The NRC report (2) was very clear on this point and presented a quite lengthy discussion on the nature of and magnitude of errors, which the Asbury et al. report does not consider. For example, contrary to the assertion of Asbury et al., "... in favor of direct comparison of alkalinity values which can be determined with great precision ..." (emphasis mine), the methyl orange
0 1990 American Chemical Society
Environ. Sci. Technol., Vol. 24, No. 3, 1990 385
