Aquatic chemistry of acid deposition Several chemical concepts are prerequisite to understanding acid deposition genesis and properties
Werner Stumm Laura S L
Swiis Federal Institute of Technology Zurich. Switzerland Jerald L. Schnoor University of l o w lowa City, Iowa 52242
The Occurrence of acid precipitation in many regions of the Northern Hemisphere is a consequence of human interference in the cycles that unite land, water, and atmosphere. The oxidation of carbon, sulfur, and nitrogen, resulting mostly from fossil fuel burning, rivals oxidation processes induced by photosynthesis and respiration and disturbs redox conditions in the atmosphere. The atmosphere is more suscep tible to anthropogenic emissions than are the terrestrial or aqueous environments because, from a quantitative point of view, the atmosphere is much smaller than the other reservoirs are. Furthermore, the time constants concerning atmospheric alterations are small in comparison with those of the seas and the lithosphere (I, 2). In oxidation-reduction reactions, electron transfers (e-) are coupled with the transfer of protons (H+)to maintain a charge balance. A modification of the redox balance corresponds to a moditication of the acid-base balance (3).The net reactions of the oxidation of C, S, and N exceed reduction reactions in these elemental cycles. A net production of Hf ions in atmospheric precipi8 Environ. Sci. Technol.. MI.21. NO.1. 1987
lust, ocean aemls. HC, metals
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0013-936X18610921-008$01 5010
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1986 American Chemical Society
tation is a necessary consequence. The disturbance is transferred to the terrestrial and aquatic environments, and it can impair terrestrial and aquatic ecosystems. Figure 1 shows the various reactions that involve atmospheric pollutants and natural components in the atmosphere. The following reactions are of particular importance in the formation of acid precipitation: oxidative reactions, either in the gaseous phase or in the aqueous phase, leading to the formation of oxides of C, S, and N (C02; SO2, SO,, H2S04; NO, NO,, "02, "01); absorption of gases into water (cloud droplets, falling raindrops, or fog) and interaction of the resulting acids (SQ.H20, H2S04,"0,) with ammonia ("3) and the carbonates of airborne dust; and the scavenging and partial dissolution of aerosols into water. In this
1
case, aerosols are produced from the
interaction of vapors and airborne (maritime and dust) particles; they often contain (NH4)2S04 and NRNO,. The products of the various chemical and physical reactions are eventually returned to the Earth's surface. Usually, one distinguishes between wet and dry deposition. Wet deposition (rainout and washout) includes the flux of all those components that are carried to the Earth's surface by rain or snow, that is, those dissolved and particulate substances contained in rain or snow. Dry deposition is the flux of particles and and gases (especially S02, "03, NH,) to the receptor surface during the absence of rain or snow (47) Deposi. tidn also can occur through fog aerosols and droplets, which can be deposited on trees, plants, or the ground. It is beyond the scope of this article to review in any detail the various oxidation and scavenging processes that are essential to the formation of acid deposition. Instead, we concentrate on three elementary chemical concepts that are the prerequisites to understanding the genesis and properties of acid deposition, thus illustrating that the principles conventionally used in water chemistry can be applied to the chemistry of atmospheric water. First of the three concepts is a simple stoichiometric model, which explains on a mass balance basis that the composition of the rain results primarily from a titration of the acids formed from atmospheric pollutants with the bases (NH,- and C0,2-lxaring dust particles) introduced into the atmosphere. Next is an illustration of the absorption equilibria of such gases as SQ and NH, into water, which represents their interaction with cloud water, raindrops, fog droplets, or surface waters. Finally, a quantitative interpretation of the acidbase balance is presented. This requires a rigorous definition of acid-neutralizing capacity (alkalinity) and base-neutralizing capacity (acidity) to measure the residual acidity of acid precipitation and to interpret the interaction of acid deposition with the terrestrial and aquatic environment. Stoichiometric model The rain water shown in Figure I contains an excess of strong acids, most of which originate from the oxidation of sulfur during fossil fuel combustion and from the fixation of atmospheric nitrogen to NO and NO, (for example, during combustion of gasoline by mw tor vehicles). It also should be mentioned that there are natural sources of acidity, resulting from volcanic activity, from H2S from anaerobic sediments, and from dimethyl sulfide and carbonyl
sulfide that originate in the Oman. H('I results from the combustion and dccomposition of organochlorine compounds such as polyvinyl chloride. Bases originate in the atmosphere as the carbonate of wind-blown dust and from NH2, generally of natural origin. The NH, comes from N R + and from the decomposition of urea in soil and agricultural environments. The reaction rates for oxidation of atmospheric SO2 (0.05-0.5.day-') yield a sulfur residence time of several days at the most; this corresponds to a transport distance of several hundred to 1030 km (8).The formation of HNO, by oxidation is more rapid and, compared with H2S04, results in a shorter travel distance from the emission source (3). H2S04 also can react with NH, to form ",HS04 or (NH.,),S04 aerosols. In addition, the NH4N03 aerosols are in equilibrium with NH,(g) and HNO,(g)(9). Figure 1 illustrates wet and dry deposition. The flux of dry deposition is usually assumed to be a product of its concentration adjacent to the surface and the deposition velocity. Deposition velocity depends on the nature of the pollutant (type of gas, particle size), the turbulence of the atmosphere, and the characteristics of the receptor surface (water, ice, snow, vegetation, trees, rocks). Dry deposition is usually collected for study in open buckets during dry periods, but this method underestimates the flux of pollutants, particularly gases (SO,, NO,), to a foliar canopy or lake surface. (Fog can be collected with special devices, but it is not measured quantitatively in a wet or ,dry bucket collector.) The foliar canopy receives much of its dry deposition in the form of sulfate, nitrate, and hydrogen ion, which occur primarily as SO,, HNO,, and NH, vapors. Dry deposition of coarse particles has been shown to be an important source of calcium and potassium ions deposited on deciduous forests in the eastern United States (7). In addition to the acid-base components shown in Figure I , various organic acids are often found. Many of these acids are byproducts of the atmospheric oxidation of organic matter released into the atmosphere (1G13).Of special interest are formic, acetic, oxalic, and benzoic acids, which have been found in rain water in concentrations occasionally exceeding a few micromoles per liter. Thus, they must be included in the calculated acidity of rain water, and their presence in larger concentrations has an influence on the pH of rain and fog samples. Figure 2 illustrates the inorganic composition of representative rain samples (14).The ratio of the cations (H+, Environ. Sci. Technol.,Vol. 21. No. 1. 1987 9
N&+, Ca2+, Na+, and K+) and the anions NO3-, Cl-) reflects the acid-base titration that occurs in the atmosphere and in rain droplets. Total concentrations (the sum of cations or anions)typically vary from 20 MIL to 500 fieq/L. Dilution effects, such as washont by atmospheric precipitation. can in part explain the differences observed. For example, the highest concenmtions are observed after long, dry perioaS; the lowest concentrations are typically recorded at the end of prolonged rain. When fog is formed from water-sahrated air; water droplets condense on aerosol particles. In addition to components of the aerosols, the fog droplets can absorb such gases as NO,. SQ, NH,, and H a ; they form a favorable milieu for various oxidation processes, especially the formation of H2S04.Fog droplets (10-50 pm in diameter) are much smaller rain droplets are;the liquid water content of fog is often in the range of 1 x IO4Urn3 air, so that acids in fog are typically 10-50 times 10 Envimn. Scl. Technol., Vol. 21. No. 1. 1987
more concennated than are those found in rain (Figure 2) (14). Rain clouds process a considerable volume of air over relatively large distances and thus are able to absorb gases and aerosols from a large region. B e cause fog is formed in the lower air masses, fog droplets are efficient collectors of pollutants close to the Earth’s surface. The influence of local emissions (such as MI3 in agricultural regions or HCl near refuse incinerators) is reflected in the fog composition. Fog waters typically contain total ionic concentrations of 0.5-15 meq/L,as shown in Figure 2. Remarkably different pH values can be observed in fog. In addition to neutral fogs @H 5-7)-some of which have very high anion concentrations-we have found, as have others, extreme acidity @H 1.8) (15).
Sq and N H 3 adsorption The distribution of gas molecules between the gas phase and the water phase depends on the Henry’s law eqnilibrium distribution. In the case of C02,
a,
and NH,, the dissolution equilibrium is pH dependent because the components in the water phase-C02(aq), H2C4, S02’H20(aq), NH3(aq)-undergo acid-base reactions. W o varieties of calculation are possible (2, 16-18): In an open-system model, a constant partial pressure of the gas component is maintained (Figure 3a). In a closed system, an initial partial pressure of a component is given, for example, for a cloud before rain droplets are formed or for a package of air before fog droplets condense. In these cases, the system is considered closed-from then on, the total concentration in the gas phase and in the solution phase is constant (Figure 3b). These examples illustrate the strong pH dependence of the solubility of these atmospheric gases. They relate to the question of linearity of emission controls for S@. In some instances other factors may affect the solubility. For example, the presence of formaldehyde can enhance the absorption of (total) SO2 in water, because CH20 reacts
with HS03- to form CH20HS03-. Hydroxymethanesulfonic acid is a very strong acid, and oxidation of S(W) bound to it is inhibited. To calculate the equilibrium of COz, and NH3 between the gas phase and the liquid phase of atmospheric water, it is fmt necessary to define the systems and to write equilibrium equations. For equilibrium at 25 'C (infinite dilution) the C q system equations are as follows:
a,
[S4.H20]/Pw2 = 1W1[M atm-'1 (4) tHSO3-1 [H+]/[SCh.HzO] = 10-1 (5) [S032-] [H+]/@SO3-] = 10-7.2 (6) Finally, the NH3 system equations are written as follows:
These equations are plotted in Figure 3a. Similarly, for an open S q system, ["3(aq)l/P~~~ = 10' * [M am-'] (7) Psm = 2 x 10-' atm (constant), we [NH3(as)l [ H + I / W + ] = lCr9.2 (8) obtain the distribution by combining Equations 4-6 (Figure 3a). The distribution of the species in an open system (constant partial pressure) is best obtained by plotting all the perti@zC03*]/Pcoz = lW1.5[M atm-l] (1) nent equations in a double logarithmic [HCO;] [H+I/[H~CO,*I= 10-6.3 (2) diagram (log molar concentration of the individual species in water vs. pH). For (3) example, in Figure 3a we obtain ex[C0,2-] [H+]/[CO,2-] = 10-10.2 [Sa21 =
[email protected]@7.2.16.1. Where P is partial pressure and pressions for Pcq
[email protected] (composip S o z / ~ + ]=2 10-16.71[~+12(14) W Z C W = [ C 9 W ) l + [ H Z C ~ I tion of the atmosphere) by combining (4).The system equations, also Equations 1-3 to arrive at In a closed system, we can make the valid for equilibrium at 25 OC (5)(infi[H~CO~*I = pco2. 10-13 = calculation on a mass balance basis for nite dilution), are written as follows: 10-5 M (9) a given liquid water content of the at-
Sa
mosphere, q, [L m-'1. The distribution of atmospheric NH3 also is a strong function of pH. If we assume a total concentration of NH3 in gas and water of 3 x IW7 mol m-3, corresponding to an initial PNH3 (in the absence of water) of 7.3 x atm and a liquid water content of q = 5 x lW L m-3, the mass balance for total NH3 in the gas and liquid water phases is 3 x IW7 mol = NH4d
+
where RT (at 25 "C) = 2.446 x 1 W m3 a m mol-'. The partial pressure of ammonia (PNH3) and then the other species can be calculated as a function of pH (Figure 3b). The calculation can be simplified if one realizes that at high pH nearly all NH3 is in the gas phase, whereas at low pH nearly all of it is dissolved as NH, . For SO, and an assumed total concentration of 8 x IlY7 mol m-3 (an initial PSw of 2 x 1 W a m ) and a liquid water content of q = 5 x 1w4 L m-3, the overall mass balance is given by the following equation: SO, (total) = 8 x lW7 mol m-3 = Ps IRT [ P s ~ . 1 6 . ' ( 1 lO-%/[H+]
[email protected]/[H+]Z)]q (16) +
+
+
+
At low pH, nearly all SO, is in the gas phase; at high pH most of it is dissolved as S03z-(Figure 3b). Neubalizing capacity One must distinguish between the H+ concentration (or activity) as an intensity factor and the availability of H+, the H+ ion reservoir, as given by the base-neutralizing capacity, BNC, or [H-Acyl (2,3, 19,20). The BNC can be defined by a net proton balance with regard to a reference level-the sum of the concentrations of all the species containing protons in excess of the reference level, less the concentrations of the species containing protons in deficiency of the proton reference level. For natural waters, a convenient reference level (corresponding to an equivalence point in alkalimetric titrations) includes H 2 0 and H2CO3:
The acid-neutralizing capacity, ANC, or alkalinity [Alkl is related to [H-Acyl
bY -[H-Acy] = [Alk] = [HCO,-] + 2 [ C a 2 1 + [OH-] - PI+](18) Considering a charge balance for a typical natural water (Figure 4) we realizethat [Alk] and [H-Acyl also can be expressed by a charge balance-the 12 Environ. Sci. Technol., MI. 21, No. 1, I987
Naturalwaterchargebalance a
-Wequivalent sum of conservative cations, less the sum of conservative anions ([Alk]= a - b). The conservative cations are the base cations of the strong bases Ca(0H2), KOH, and the like; the conservative anions are those that are the conjugate bases of strong acids (SO,2-, NO;, and C1-). [Alk] = rnCO3-1 + 2[C032-] + [OH-] - [H+] = ma+] [K+] 2[Ca2+] 2[Mg2+]- [Cl-] 2[S0a2-] - [NO,-] (19) The [HI-Acylfor this particular water, obviously negative, is defined ([HAcyl = b - a) as IH-Awl = IC1-1 + 21SOa21+
+
+
+
The definitions given for proton balance and electroneutrality are compatible because the reference state has been defined in terms of uncharged species. These definitions can be used to interpret interactions of acid precipitation with the environment. A simple accounting can be made: Every base cationic charge unit (for example, Caz+ or K+)that is removed from the water by whatever process is equivalent to a proton added to the water, and every conservative anionic charge unit (from anions of strong acids-NO3-, SO,'-, or C1-) removed from the water corresponds to a proton removed from the water: or generallv. ZA [base cations] ZA [conservative anions] = +A[Ak] = -A[H-Acy] (21) If the water under consideration contains other acid- or baseans-ng species, the proton reference level must be extended to the other components. In addition to the species given, if a natural water contains organic molecules,such as a carboxylic &id ( H a g ) and ammonium ions. the reference level is extended to these species. One could choose the ammonium scecies m+ and Hdre as a zero-level rkference skte. In opeFation, we wish to distinguish between the acidity caused by strong acids (mineral acids and organic acids with pK < 6)-typiI
cally called mineral acidity orfree acidify, which often is nearly the same as the free-H+ concentration-and the total acidity given by the BNC of the sum of smng and weak acids (20). The distinction is possible by careful alkalimetric titrations of rain and fog samples. Gran titrations have found wide acceptance in this area. The strong acidity, [H-Acyl, is expressed with respect to the reference conditions: H20, H2C03, SO,", NO3-, CI-, NO,-, P,HS03-, W+,H,Si04, and EBOrg (organic acids, pK. < IO).
components such as HS04-, HNO,, HF, H2S03, C032-, NH,, and H3Si04are in negligible concentrations in typical rain water. Thus, the equation may be simplified to [H-Acyl = [H+] - [OH-] [HC03-] - Xn[Org"l (23) For most rain samples of pH 4-4.5, [H-Acyl is equal to p+], but in highly concentrated fog waters (in extreme cases, pH
