10. ATMOSPHERIC DEPOSITION AND BIOGEOCHEMISTRY

1. 10. ATMOSPHERIC DEPOSITION AND BIOGEOCHEMISTRY For in the end we will conserve only what we love. We will love only what we understand. And we will understand only what we are taught. Baba Dioum, African Conservationist

2. Atmosphere, land, and water are interconnected compartments. It makes no sense to speak solely of water pollution because of intermedia transfers to air or land. • If one removes pollutions from a wastewater discharge in order to improve water quality, residuals are deposited onto land or, if incinerated, into the air. • The atmosphere transfers pollutants to land and water via atmospheric deposition, that is, the transport of pollutants, both gaseous and particulate, from the air to land and water. • Acid precipitation is the most common form of atmospheric deposition, and it affects elemental cycling in the environment (biogeochemistry) and heavy metals transport.

3. 10.1 GENESIS OF ACID DEPOSITION • The oxidation of carbon, sulfur, and nitrogen, resulting from fossil fuel burning, disturbs redox conditions in the atmosphere. The atmosphere is more susceptible 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. • 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 modification of the acid-base balance. • Figure 10.1 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 or C, S and N (CO2; SO2, SO3, H2SO4; NO, NO2, HNO2, HNO3); absorption of gases into water (cloud droplets, falling raindrops, or fog) and interaction of the resulting acids (SO2 · H2O, H2SO4, HNO3) with ammonia (NH3) and the carbonates of airborne dust, and the scavenging and partial dissolution of aerosols into water.

4. Figure 10.1 Depiction of the genesis of acid rain. From the oxidation of S and N during the combustion of fossil fuels, there is a buildup in the atmosphere (in the gas phase, aerosol particles, raindrops, snowflakes, and fog) of CO2 and the oxides of S and N, which leads to acid-base interaction. The importance of absorption of gases into the various phases of gas, aerosol, and atmospheric water depends on a number of factors. The genesis of acid rain is shown on the upper right as an acid-base titration. Various interactions with the terrestrial and aquatic environment are shown in the lower part of the figure.

5. 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 gases (especially SO2, HNO3, and NH3) to the receptor surface during the absence of rain or snow. • Three elementary chemical concepts are prerequisites to understanding the genesis and modeling of acid deposition. 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 (NH3- and CO32- - bearing dust particles) introduced into the atmosphere. • Next is an illustration of the absorption equilibria of such gases as SO2 and NH3 into water, which represents their interaction with cloud water, raindrops, fog droplets, or surface waters.

6. 10.1.1 Stoichiometric Model • The rainwater shown in Figure 10.1 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 NO2 (e.g., during combustion of gasoline by motor vehicles). In addition, 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 ocean. • Reaction rates for the oxidation of atmospheric SO2 (0.05-0.5 day-1) yield a sulfur residence time of several days at the most; this corresponds to a transport distance of several hundred to 1000 km. The formation of HNO3 by oxidation is more rapid and, compared with H2SO4, results in a shorter travel distance from the emission source. H2SO4 also can react with NH3 to form NH4HSO4 or (NH4)2SO4 aerosols. • 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 atmosphere, and the characteristics of the receptor surface (water, ice, snow, vegetation, trees, rocks).

7. The foliar canopy receives much of its dry deposition in the form of sulfate, nitrate, and hydrogen ions, which occur primarily as SO2, HNO3, and NH3 vapors. Dry deposition of coarse particles has been shown to be an important source of calcium and potassium ion deposition on deciduous forests in the eastern United States (Table 10.1). • Figure 10.1 show the acid-base components. Many of these acids are by-products of the atmospheric oxidation of organic matter released into the atmosphere. Of special interest are formic, acetic, oxalic, and benzoic acids, which have been found in rainwater in concentrations occasionally exceeding a few micromoles per liter. • Figure 10.2 illustrates the inorganic composition of representative rain samples. The ratio of the cations (H+, NH4+, Ca2+, Na+, and K+) and the anions (SO42-, 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 µeq L-1 to 500 µeq L-1. Dilution effects, such as washout by atmospheric precipitation, can in part explain the differences observed. • When fog is formed from water-saturated air, water droplets condense on aerosol particles. • Typical water contents in atmospheric systems are 5 × 10-5 to 5 × 10-3 L m-3 for fog and 10-4 to 10-3 L m-3 for clouds.

8. Table 10.1 Total Annual Atmospheric Deposition of Major Ions to an Oak Forest at Walker Branch Watershed, Tennessee.

9. Figure 10.2 Composition of fog and rain samples, in a highly settled region around Zurich, Switzerland. The composition of fog varies widely and reflects to a larger extent than rain the influence of local emissions close to the ground. The fog concentration increases with decreasing liquid water content.

10. 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. Because 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 NH3 in agricultural regions or HCl near refuse incinerators) is reflected in the fog composition. • Table 10.2 is a summary of KH (Henry's constant and other equilibrium constants at 25ºC (for Henry's law Caq = KH patm) for most gases of importance in atmospheric deposition to lakes and forests. Henry's law constants, as for other thermodynamic constants, are valid for ideal solutions. • Ideally, they should be written in terms of activities and fugacities. Since activity coefficients for neutral molecules in aqueous solution become larger than 1.0 (salting-out effect), the solubility of gases is smaller in salt solution than in dilute aqueous medium (expressed in concentration units).

11. Table 10.2 Equilibrium Constants of Importance in Fog-Water Equilibrium

12. Example 10.1 Solubility of SO2 in Water • a. What is the solubility of SO2 (Patm = 2×10-9 atm) in water at 5 ºC? • b. What is the solubility of CO2 (Patm = 3.3×10-4 atm) in water at 5 ºC? • c. What is the composition of rain in equilibrium with both SO2 and CO2 in water under the conditions specified? Solution: • a. The following constants are valid at 5ºC after correction using the van‘t Hoff relationship:

13. The solubility of SO2 can be calculated the same way as that of CO2. The calculation for SO2 solubility is as a function of pH. If SO2 alone (no acids or bases added) comes into contact with water droplets, the composition is given approximately at a pH where [H+] = [HSO3-] (see Figure 10.3a). The exact proton condition or charge balance condition is TOT H = (H+) - (HSO3-) - 2(SO32-) – [OH-] = 0. This condition applies to the following composition: pH = 4.9, [HSO3-] = 1.2 × 10-6 M, [SO2· H2O] = 3.7× 10-8 M, [SO32-] = 8 × 10-8 M. • b. For this part we have pH = 5.65, [H2CO3*] = 2.1 × 10-5 M, [HCO3-] = 2.2 × 10-6 M, [CO32-] = 2.7 × 10-11 M. The answer to question (b) is obtained by plotting the corresponding diagram for CO2 where [H+] = [HCO3-]. • c. The answer to question (c) is obtained by superimposing the plots for SO2 and for CO2. The matrix for solution of the chemical equilibrium problem is given by Table 10.3.

14. Electroneutrality [equation (ix)] specifies the condition for water in equilibrium with the given partial pressures of CO2 and SO2 (no acid or base added). This condition is fulfilled where [H+] ≈ [HSO3-] + [HCO3-] + 2[SO32-]. • SO2, even at small concentrations, has an influence on the pH of the water droplets. It has shifted from pH = 5.6 (where [H+] = [HCO3-]) to pH = 4.75. The exact answer for this composition is in logarithmic units (Table 10.4). • The effect of CO2 on the pH of the system is very small compared to that of SO2. • The units for Henry`s law constants in Table 10.2 are expressed as M atm-1, but oftentimes they are given in inverse units in the literature, so one must be careful. • Here, we will use KH in M atm-1 and R = 0.08206 atm M-1 K-1 (at 25 ºC, RT = 24.5 atm M-1).

15. Table 10.4 Equilibrium Compositiona for Example 10.1c

16. 10.1.2 SO2 and NH3 Absorption • The distribution of gas molecules between the gas phase and the water phase depends on the Henry‘s law equilibrium distribution. In the case of CO2, SO2, and NH3, the dissolution equilibrium is pH dependent because the components in the water phase - CO2(aq), H2CO3, SO2· H2O(aq), NH3(aq) - undergo acid-base reactions. • Two varieties of chemical equilibrium modeling are possible. In an open system model, a constant partial pressure of the gas component is maintained. 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 this case, the system is considered closed from then on, the total concentration in the gas phase and in the solution phase is constant (Figure 10.3). • For equilibrium at 25ºC (infinite dilution) the CO2 system equation are as follows: (1) (2) (3)

17. Where P is partial pressure and [H2CO3*] = [CO2 · (aq)] + [H2CO3]. The SO2 system equations, also valid for equilibrium at 25ºC (infinite dilution), are written as follows: (4) (5) (6) • Finally, the NH3 system equations are written as follows: (7) (8)

18. For example, in Figure 10.3a we obtain expressions for PCO2 = 10-3.5 atm (composition of the atmosphere) by combining equations (1)-(3) to arrive at (9) (10) (11) • These equations are plotted in Figure 10.3a. Similarly, for an open SO2 system, PSO2 = 2 × 10-8 atm (constant), we obtain the distribution by combining equations (4)-(6) (Figure 10.3a). (12) (13) (14) • The closed system model can often be used expeditiously when a predominant fraction of the species is absorbed in the water phase. In a closed system, the total concentration is constant.

19. Figure 10.3 Equilibria with the atmosphere (atmospheric water droplets) for the conditions given. • Open systems: atmospheric CO2 with water, PCO2 = 10-3.5; PSO2 = 2 ×10-8. • Closed systems: atmospheric NH3 with water, liquid water content 5 ×10-4 L m-3; total NH3 = 3 × 10-7 mol m-3; total SO2 = 8 × 10-7 mol m-3.

20. The mass balance for total NH3 in the gas and liquid water phases is (15) • where RT (at 25 ºC) = 2.446 × 10-2 m3 atm mol-1. The partial pressure of ammonia (PNH3) and then the other species can be calculated as a function of pH (Figure 10.3b). 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 NH4+. At low pH, water vapor is an efficient sorbent for NH3 gas, but it decreases at higher pH. • For SO2 and an assumed total concentration of 8 × 10-7 mol m-3 (an initial PSO2 of 2 × 10-2 atm) and a liquid water content of q = 5 × 10-4 L m-3, the overall mass balance is given by the following equation: (16)

21. 10.2 ACIDITY AND ALKALINITY; NEUTRALIZING CAPACITIES • One has to distinguish between the H+ concentration (or activity) as an intensity factor and the availability of H+, that is, the H+ -ion reservoir as given by the base-neutralizing capacity, BNC. The BNC relates to the alkalinity [Alk] or acid-neutralizing capacity, ANC, by (17) • For natural waters, a convenient reference level (corresponding to an equivalence point in alkalimetric titrations) includes H2O and H2CO3: (18) • The acid-neutralizing capacity, ANC, or alkalinity [Alk] is related to [H-Acy] by (19) • Considering a charge balance for a typical natural water (Figure l0.4), we realize that [Alk] and [H-Acy] also can be expressed by a charge balance: the equivalent sum of conservative cations, less the sum of conservative anions ([Alk] = a – b).

22. Figure 10.4 Natural water charge balance for an alkaline system (Alk = a - b) and an acid system (Alk = a – b = d – c)

23. The conservative cations are the base cations of the strong bases Ca(OH)2, KOH, and the like; the conservative anions are those that are the conjugate bases of strong acids (SO42-, NO3- and Cl-). (20) • The [H-Acy] for this particular water, obviously negative, is defined ([H-Acy] = b - a) as (21) • These definitions can be used to interpret interaction of acid precipitation with the environment. • A simple accounting can be made: (22) • If the water under consideration contains other acid- or base-consuming species, the proton reference level must be extended to the other components.

24. In operation, we wish to distinguish between the acidity caused by strong acids (mineral acids and organic acids with pK < 6) typically called mineral acidity or free acidity, which often is nearly the same as the free-H+ concentration, and the total acidity given by the BNC of the sum of strong and weak acids. • The distinction is possible by careful alkalimetric titrations of rain and fog samples. Gran titrations have found wide acceptance in this area. (23) • Components such as HSO4-, HNO2, HF, H2SO3, CO32-, NH3, and H3SiO4 are in negligible concentrations in typical rainwater. Thus the equation may be simplified to (24) • For most rain samples of pH 4-4.5, [H-Acy] is equal to [H+], but in highly concentrated fog waters (in extreme cases, pH < 2.5) HSO4- and SO2· H2O become important species contributing to the strong acidity.

25. The reference conditions pertaining to the determination of total acidity [AcyT] are H2O, CO32-, SO42-, NO3-, Cl-, NO2-, F-, SO32-, NH3, H3SiO4-, ΣOrgn-, and Al(OH)3. (25) • For most sample this equation can be simplified as (26) • Gran titration of the strong acidity usually gives a good approximation of the acidity [H-ACy], as defined above, but one must be aware that organic acids with pKa 3.5-5 are partly included in this titration and may affect the resulting Gran functions.

26. 10.2.1 Atmospheric Acidity and Alkalinity • In a (hypothetically closed) large system of the environment consisting of the reservoir atmosphere, hydrosphere, and lithosphere, a proton and electron balance is maintained. Temporal and spatial inhomogeneities between and within individual reservoirs cause significant shifts in electron and proton balance, so that subsystems contain differences in acidity or alkalinity. Any transfer of an oxidant or reductant, of an acid or base, or of ions from one system to another (however caused, by transport, chemical reaction, or redox process) causes a corresponding transfer of acidity or alkalinity. • Morgan, Liljestrand, and Jacob et al. introduced the concept of atmospheric acidity and alkalinity to interpret the interactions of NH3 with strong acids emitted into and/or produced within the atmosphere. • Figure 10.5 exemplifies the concept of alkalinity, Alk, and acidity, Acy, for a gas-water environment and defines the relevant reference conditions. • In Figures 10.5a and 10.5b, it is shown how the gases NH3, SO2, NOx, HNO3, HCl, and CO2 (potential bases or acids, respectively), subsequent to their dissolution in water and the oxidation of SO2 to H2SO4 and of NOx to HNO3, become alkalinity or acidity components.

27. Figure 10.5 Alkalinity/acidity in atmosphere, aerosols, and atmospheric water. Alkalinity and acidity can be defined for the atmosphere using a reference state valid for oxide conditions (SO2 and NOx oxidized to H2SO4 and HNO3) and in the presence of water. The neutralization of atmospheric acidity by NH3 is a major driving force in atmospheric deposition.

28. Thus {Alk}(gas) and {Acy}(gas), for the gas phase, is defined by the following relation: (27) and (28) • where { } indicate mol m-3 and H-Org is the sum of volatile organic acids. Figure 10.5b shows that these potential acids and bases, subsequent to their dissolution in cloud of fog water with q = 10-4 L H2O per m3 atmosphere, give a water with the equivalent acidity. • In the case of aerosols, we can define the alkalinity by a charge balance of the sum of conservative cations, Σn{Cat.n+}(ae), of NH4+, {NH4+}(ae), and of the sum of conservative anions, Σm{An.m-}(ae) (Figure 10.5c): (29)

29. At low buffer intensity, for example, in the case of residual atmospheric acidity production alleviates further SO2 oxidation, if (30) • Thus, while NH3 introduced into the atmosphere reduces its acidity, it enhances the oxidation of SO2 by ozone, participates in the formation of ammonium sulfate and ammonium nitrate aerosols, and accelerates the deposition of SO42-. Furthermore, any NH3 or NH4+ that is returned to the earth's surface becomes HNO3 as a consequence of nitrification and/or H+ ions as a consequence of plant uptake, (31) • and it may aggravate the acidification of soils and lakes. This effect is not sufficiently considered in the assessment of NH3 emissions (e.g., agriculture, feed lots) and the use of excess NH3 in air pollution control processes to reduce nitrogen oxides.

30. Example 10.2 Mixing of Water with Different Acid-Neutralizing Capacities • The effluent from an acid lake with [H-Acy] = 5 × 10-5 eq L-1 and a pH of 4.3 mixes with a river containing an [Alk] = 1.5 × 10-4 eq L-1 and a pH of 7.4 in a 1:1 volumetric ratio. What is the alkalinity and the pH of the mixed waters? You may assume that the mixed water is in equilibrium with the CO2 of the atmosphere (3.5 × 10-4 atm) and at 10ºC. The acidity constant of H2CO3* is 3 × 10-7 and Henry's constant for the reaction CO2(g) + H2O = H2CO3* is KH = 0.050 M atm-1. • Solution: Alkalinity (ANC) is a conservative quantity that is unaffected by CO2(g) sorption. We may calculate the alkalinity of the mixture by volume-weighted averaging.

31. The concentration of H2CO3* in equilibrium with the atmosphere is given by • At a pH of 7.4, most of the alkalinity is due to HCO3- so that [Alk] = [HCO3-] = 10-4 M. Then, [H+] is given by the equilibrium expression at 10ºC: • and [H+] = 5.2 × 10-8; pH = 7.3, close to that of the original river. • This example illustrates that (1) [Alk] = -[H-Acy], (2) [Alk] and [H-Acy] are conservative parameters and can be used directly in mixing calculations, and (3) [H+] and pH are not conservative parameters. The river was well buffered by the bicarbonate system despite an equal volume of acid input at low concentration.

32. 10.3 WET AND DRY DEPOSITION 10.3.1 Wet Deposition • Wet deposition occurs when pollutants fall to the ground or sea by rainfall, snowfall, or hail/sleet. Dry deposition is when gases and aerosol particles are intercepted by the earth's surface in the absence of precipitation. Let us first discuss wet deposition. Wet deposition to the surface of the earth is directly proportional to the concentration of pollutant in the rain, snow, or ice phase. • The wet deposition flux is defined by equation (32): (32) • where Fwet, is the areal wet deposition flux in µg cm-2s-1, I is the precipitation rate in cm s-1 (as liquid H2O), and Cw is the concentration of the pollutant associated with the precipitation in µg cm-3. • The concentration of pollutants in wet deposition is due to two important effects with quite different physical mechanisms: • Aerosol particle scavenging. • Gas scavenging

33. Aerosols begin their life cycle after nucleation and formation of a submicron hydroscopic particle, for example, (NH4)2 SO4, which hydrates and grows very quickly due to condensation of water around the particle. At this stage, it is neither solid nor liquid, but merely a stable aerosol with a density between 1.0 and 1.1 g cm-3. • Assuming an average spacing of 1-mm between cloud droplets, condensation of 106 cloud droplets into a 1-mm raindrop would scavenge enough air for a washout ratio of 106 (33) • where Cw is the concentration of the pollutant in precipitation water in µg cm-3, Cae is the concentration of the pollutant associated with aerosol droplets in air in µg cm-3, and W is the washout ratio for aerosols, dimensionless (cm3 air/ cm3 precipitation). • Table 10.5 provides a few values of washout ratios for metals associated with particles, and they are typically on the order of 105-106. Rainout sometimes refers to below-cloud processes, whereby pollutants are scavenged as raindrops fall through polluted air.

34. Table 10.5 Some Measured Values for Size and Washout Ratio of Metals as Aerosols in the Atmosphere

35. If we express Henry's constant KH in units of M atm-1, the following equations apply for Henry's law and the washout ratio: (34) (35) • where Cw is the concentration in the water phase (M), patm is the atmospheric partial pressure (atm), W is the washout ratio (dimensionless, i.e., L H2O/L gas), Cg is the concentration in the gas (mol L-1 gas), and RT is the universal gas law constant times temperature (24.46 atm M-1 at 25ºC). • Table 10.6: some estimates for washout ratios of selected pesticides. Henry‘s constants are taken from Schwarzenbach et al.. In general, washout ratios are large for soluble and polar compounds, intermediate for semivolatile chemical (such as DDT, dieldrin, dioxin, and PCBs), and low for volatile organic chemicals. Semivolatile pollutants are an interesting case because these gases can be transported long distances and recycled many times before being deposited in polar regions by a "cold-trap" effect.

36. Table 10.6 Estimates of Washout Ratios for Selected Gases, 25ºC

37. Example 10.3 Washout of Pollutants from the Atmosphere • To what extent are atmospheric pollutants washed out by rain? We can try to answer this question by considering the gas absorption equilibria. Our estimate is based on the following assumptions and mass balance considerations. For example, calculate the mass fraction that is washed out (fwater) for the pesticide lindane (γ-hexachlorocyclohexane, C6H6Cl6) with Henry’s constant KH of 309 M atm-1. • Solution: Assume the height of the air column is 5 × 103 m. This column is “washed out” by a rain of 25 mm (corresponding to 25 L m-2). In other words, gas volume Vg = 5 × 103 m3 water volume Vw = 0.025 m3 • The total quantity of the pollutant is

38. The fraction of pollutants in the water phase, fwater, is given by • Lindane is quite soluble, relatively speaking, but only about 3.65% of it is washed out by the rainfall.

39. 10.3.2 Dry Deposition • Both wet and dry deposition are important transport mechanisms. For total sulfur deposition in the United States, they are roughly of equal magnitude. Dry deposition takes place (in the absence of rain) by two pathways: • Aerosol and particle deposition. • Gas deposition. • There are three resistances to aerosol and gas deposition: (1) aerodynamic resistance, (2) boundary layer resistance, and (3) surface resistance. • Aerodynamic resistance involves turbulent mixing and transport from the atmosphere (~1-km elevation) to the laminar boundary layer in the quiescent zone above the earth's surface. • Dry deposition velocity encompasses the electrical analog of these three resistance in series: (36) • where Vd is defined as the dry deposition velocity (cm s-1), ra is the aerodynamic resistance, rb is the boundary layer resistance, and rs is the resistance at the surface.

40. The deposition velocity is affected by a number of factors including relative humidity, type of aerosol or gas, aerosol particle size, wind velocity profile, type of surface receptor, roughness factor, atmospheric stability, and temperature. Vd increases with wind speed because sheer stress at the surface causes increased vertical turbulence and eddies. • For aerosol particles, the deposition velocity is dependent on particle diameter as shown in Figure 10.6. Milford and Davidson showed a general power-law correlation for the dependence of Vd on particle size: (37) • where Vd is the deposition velocity in cm s-1 and MMD is the mass median diameter of the particle in µm. • Table 10.7 is a compilation of dry deposition velocities for chemicals of interest from Davidson and Wu.

41. Figure 10.6 Dry deposition velocity as a function of particle diameter. Deposition velocity is always greater than the Stokes law discrete particle settling velocity (Vg) because of turbulent mixing and reaction at the surface. For very fine aerosols (less than 0.1 µm), the curve follows mass transfer correlations of the Schmidt number Sc-2/3.

42. Table 10.7 Dry Deposition Velocities for a Number of Aerosol Particles and Gases

43. In general, eases that react at the surface (e.g., SO2, HNO3, HCl, and O3) tend to have slightly higher deposition velocities, on the order of 1.0 cm s-1. HNO3 vapor has a very large deposition velocity because there is no surface resistance - it is immediately absorbed and neutralized by vegetation and/or water. • Deposition velocities in Table 10.7 are mostly to natural earth surfaces. Natural vegetation and trees are relatively efficient interceptors of gases and particles based on specific surface areas. SO2 dry deposition velocity for a coniferous forest may be several times higher than for an open field or a snow field. • Metals associated with wind-blown dust and coarse particles (Ca, Mg, K, F, Mn) tend to have higher deposition velocities due to the effect of particle size. (38) • where Xair and Alair represent the airborne concentrations of any element X and aluminum, respectively, and Xcrust and Alcrust, are the concentrations in the earth's crust. • Ag, As, Cd, Cu, Zn, Pb, and Ni tend to be enriched relative to aluminum, indicating anthropogenic origin in the atmosphere.

44. 10.4 PROSSESES THAT MODIFY THE ANC OF SOILS AND WATERS 10.4.1 ANC of Soil • In weathering reactions, alkalinity is added from the soil-rock system to the water: (39) • The acid-neutralizing capacity of a soil is given by the bases, carbonates, silicates, and oxides of the soil system. • If the composition of the soil is not known but its elemental analysis is given in oxide components, the following kind of accounting is equivalent to that given by equation (20) for natural waters: (40) • Equation (40) is expressed in oxide equivalents of each element in soil. Sulfates, nitrates, and chlorides incorporated or adsorbed are subtracted from ANC.

45. Table 10.8 Some Processes that Modify the H+ Balance in Waters

46. 10.4.2 Chemical Weathering • Figure 10.7 shows some processes that affect the acid-neutralizing capacity of soils. Ion exchange occurs at the surface of clays and organic humus in various soil horizons. The net effect of ion exchange processes is identical to chemical weathering (and alkalinity); that is, hydrogen ions are consumed and basic cations (Ca2+, Mg2+, Na+, K+) are released. • However, the kinetics of ion exchange are rapid relative to those of chemical weathering (taking minutes compared to hours or even days). In addition, the pool of exchangeable bases is small compared to the total ANC of the soil [equation (40)]. • Thus there exists two pools of bases in soils – a small pool of exchangeable bases with relatively rapid kinetics and a large pool of mineral bases with the slow kinetics of chemical weathering. • In the long run, chemical weathering is the rate-limiting step in the supply of basic cations for export from watersheds. The chemistry of natural waters is predominantly kinetically controlled.

47. Figure 10.7 Processes affecting the acid-neutralizing capacity of soils (including the exchangeable bases, cation exchange, and mineral bases). H+ ions from acid precipitation and from release by the roots react by weathering carbonates, aluminum silicates, and oxides and by surface complexation and ion exchange on clays and humus. Mechanical weathering resupplies weatherable minerals. Lines drawn out indicate flux of protons; dashed lines show flux of base cations (alkalinity). The trees (plants) act like a base pump.

48. There are several factors that affect the rate of chemical weathering in soil solution. These include: • Hydrogen ion activity of the solution • Ligand activities in solution • Dissolved CO2 activity in solution • Temperature of the soil solution • Mineralogy of the soil • Flowrate through the soil • Grain size of the soil particles. • For a given silicate mineral, the hydrogen ion activity contributes to the formation of surface-activated complexes, which determine the rate of mineral dissolution at pH < 6. • Also, since chemical weathering is a surface reaction-controlled phenomenon, organic and inorganic ligands (e.g., oxalate, formate, succinate, humic and fulvic acids, fluoride, and sulfate ) may form other surface-activated complexes that enhance dissolution. • Dissolved carbon dioxide accelerates chemical weathering presumably due to its effect on soil pH and the aggression of H2CO3.

49. Mineralogy is of prime importance. The Goldich dissolution series, which is roughly the reverse of the Bowen crystallization series, indicates that chemical weathering rates should decrease as we go from carbonates → olivine → pyroxenes → Ca, Na plagioclase → amphiboles → K-feldspars → muscovite → quartz. • If the pH of the soil solution is low enough (pH < 4.5), aluminum oxides provide some measure of neutralization to the aqueous phase along with an input of monomeric inorganic aluminum. • The key parameter that affects the activated complex and determines the rate of dissolution of the silicates is the Si/O ratio. • The less the ratio of Si/O is, the greater its chemical weathering rate is. Anorthite and forsterite have Si/O of 1:4, while quartz has Si/O of 1:2, the slowest to dissolve in acid.

50. The general rate law may be expressed as (41) • where R is the proton or ligand-promoted dissolution rate (mol m-2 s-1), k is the rate constant (s-1), Xa denotes the mole fraction of dissolution active sites (dimensionless), Pj represents the probability of finding a specific site in the coordinative arrangement of the activated precursor complex, and Ssites, is the total surface concentration of sites (mol m-2). • The rate expression in equation (41) is essentially a first-order reaction in the concentration of activated surface complex, Cj (mol m-2): (42) • Formulation of equation (42) is consistent with transition state theory, where the rate of the reaction far from equilibrium depends solely on the activity of the activated transition state complex.