Fundamentals of air Pollution – Acid Precipitation Yaacov Mamane Visiting Scientist NCR, Rome Dec May 2007 CNR, Monterotondo, Italy

Sandstone portal Figure on Herten Castle in Ruhr district of Germany, Sculpted photographed in 1908 photographed in 1969

F U M I F U G I U M: or The Inconveniencie of the AER AND SMOAK of LONDON DISSIPATED. By John Evelyn, 1661 It is this horrid Smoake which obscures our Churches, and makes our Palaces look old, which fouls our Clothes, and corrupts the waters, so as the very Rain, and refreshing Dews which fall in the several Seasons, precipitate this impure vapour, which, with its black and tenacious quality, spots and contaminates whatsoever is expos'd to it: It is this which scatters and strews about those black and smutty Atomes upon all things where it comes, insinuating it self into our very secret Cabinets, and most precious Repositories: I propose therefore, that by an Act of this present Parliament, this infernal Nuisance be reformed; enjoyning, that all those Works be removed five or six miles distant from London below the River of Thames; I say, five or six miles, or at the least so far as to stand behind that Promontory jetting out, and securing Greenwich from the pestilent Aer of Plumstead- Marshes: because, being placed at any lesser Interval beneath the City, it would not only prodigiously infect that his Majesties Royal Seat

pH Levels in the USA, 1999

History As early as 1852, R. A. Smith analyzed rain that near the industrial city of Manchester, England and found that urban aerosol particles tend to be composed primarily of sulfuric acid, but as the air is transported away from sources over more rural areas, the acid is neutralized by absorption of ammonia. urban → suburban → rural H ₂ SO ₄ + NH ₃ → (NH ₄ )HSO ₄ (+NH ₃ ) → (NH ₄ ) ₂ SO ₄ sulfuric acid → ammonium bisulfate → ammonium sulfate Throughout the early part of the twentieth century, European scientists documented the sources and effects of atmospheric acids. It was not until 1958 that acidity of precipitation in the US was characterized (Junge and Werby, 1958)

Effects - Soils Soils have colloidal molecules (clay particles) that have a layer of negative charge. They hold positively charged cations such as Al³ ⁺, K ⁺, Mg² ⁺, and Ca² ⁺. K ⁺, Mg² ⁺, Ca² ⁺ are essential plant nutrients while Al³ ⁺ is toxic. Hydrogen ions from acid deposition replace these cations on the outer layer of colloidal molecules. The metal ions are then dissolved and leached into solution and can be washed away from the soil and into surface or ground water. Soil fertilitiy is reduced and aluminum ions can replace calcium in the fish’s gills. The impact of acids on soil fertility depends on the structure and composition of the clays in the soil. The surface of the US Midwest is predominantly limestone (CaCO ₃ ), and lakes and streams have high neutralizing capacity. In the East granite dominates; soils and surface waters lacking buffering capacity, are highly sensitive to acidification.

Forests can be especially sensitive to nutrient loss. In Europe in 1993 about a quarter of the trees have died or are more than 25% defoliated. This “forest death” has been attributed at, least in part, to environmental degradation from a combination of acid deposition, ground-level ozone, and excess nutrification, primarily nitrogen. In the US, loss of forests has been so dramatic, although several species including ash and oak are sensitive to acidification of soils. Lakes and Streams The sensitivity surface waters depends critically on their neutralizing or buffering capacity. Alkaline materials such as CaCO ₃, and MgCO ₃ can neutralize acids.

Materials The Taj Mahal, the Parthenon, the Madonna in Herten, Germany, and the Lincoln Memorial are made of marble. Marble, a particular crystalline form of calcite (CaCO ₃ ), and sandstone, are subject to attack by sulfuric acid. CaSO ₄ is gypsum, which is 100 times more soluble than CaCO ₃. Many priceless historic structures have been lost to acid deposition. On a more pragmatic note, the rate of corrosion of galvanized (zinc coated) steel is 0.62 um/yr in the Adirondacks, 1.01 in Washington, DC, and 1.47 in Stubenville, OH.

Origins Primarily power generation and ore smelting. For example nickel is mined as nickel sulfide, NiS. In smelting, it is heated in air (Sudbury, Canada). The molecular weight of nickel is 57 g/mole, so smelting produces more than a ton of SO ₂ for each ton of nickel produced. Formation and Composition Gas Phase production of nitric acid: OH + NO ₂ + M → HNO ₃ + M Aqueous phase production of nitric acid: NO ₂ + O ₃ → NO ₃ + O ₂ NO ₃ + NO ₂ + M = N ₂ O ₅ + M N ₂ O ₅ + H ₂ O(l) → 2HNO ₃ (aq)

This process is important only at night, and when air temperatures are low because the formation of N ₂ O ₅ is reversible, and the equilibrium coefficient is highly temperature dependent. Also, NO ₃ is rapidly photolyzed by visible radiation. NO ₃ + hv → NO ₂ + O Gas Phase production of sulfuric acid: OH + SO ₂ + M → HOSO ₂ + M Aqueous phase production of sulfuric acid: SO ₂ + H ₂ O ₂ → H ₂ SO ₄

Global Sulfur Budget (Flux Terms In Tg S Yr -1 ) Phytoplankton (CH 3 ) 2 S SO 2  1.3d (DMS)  1.0d OHNO 3 Volcanoes Combustion Smelters SO4 2-  3.9d OH cloud dep 27 dry 20 wet dep 6 dry 44 wet H 2 SO 4 (g)

Global Sulfur Emission To The Atmosphere 1990 annual mean Chin et al. [2000]

Trends In Sulfate And Nitrate Wet Deposition

Rain Chemistry in the East Mediterranean Trends of H+ and SO4=,  eq/l

Rain Chemistry in the East Mediterranean Anions and Cations,  eq/l

AQUEOUS-PHASE CHEMISTRY HENRY’S LAW The mass of a gas that dissolves in a given amount of liquid as a given temperature is directly proportional to the partial pressure of the gas above the liquid. This law does not apply to gases that react with the liquid or ionized in the liquid. GASHENRY’S LAW CONSTANT (M / atm at 298 K) CO ₂ 3.1 x 10 ⁻ ² SO ₂ 1.3 HNO ₃ 2.1 x 10 ⁺⁵ H ₂ O ₂ 9.7 x 10 ⁺⁴

Use of Henry’s Law Assume that the atmosphere contains only N2, O2, and CO2 and that rain is in equilibrium with CO2. CO2 form a weak acid H2CO3, and it is in equilibrium with it. We should remember that: H2O = H ⁺ + OH ⁻ [H ⁺ ][OH ⁻ ] = 1 x 10 ⁻ 14 pH = -log [H ⁺ ] In pure H2O, pH = 7.0 We can assume that [CO2] in the atmosphere is around 350 ppm. ca. 370 ppm

10 -6 x350 =

K 1C = 4.3  mole/l K 2C = 4.7  mole/l K HC = M/atm

pH = 5.6

(CO2) Total = (H2CO3) + (HCO3-) + (CO3=)

K1S = 1.3  10-2 M K2S = 6.6  10-8 M KHS = 1.23 M/atm

What would be the pH of pure rain water in Rome today? Assume that the atmosphere contains only N2, O2, and CO2 and that rain is in equilibrium with CO2. Remember: H2O = H ⁺ + OH ⁻ [H ⁺ ][OH ⁻ ] = 1 x 10 ⁻ 14 pH = -log [H ⁺ ] In pure H2O, pH = 7.0 We can assume that: [CO2] = ca. 370 ppm

Today’s barometric pressure is 993 hPa = 993/1013 atm = 0.98 atm. Thus the partial pressure of CO ₂ is In water CO ₂ reacts slightly, but [H ₂ CO ₃ ] remains constant as long as the partial pressure of CO ₂ remains constant.

We know that: and Thus H + = 2.3x10 -6 → pH = -log(2.3x10 -6 ) = 5.6 EXAMPLE 2 If fog water contains enough nitric acid (HNO ₃ ) to have a pH of 4.7, can any appreciable amount nitric acid vapor return to the atmosphere? Another way to ask this question is to ask what partial pressure of HNO ₃ is in equilibrium with typical “acid rain” i.e. water at pH 4.7? We will have to assume that HNO ₃ is 50% ionized.

This is equivalent to 90 ppt, a small amount for a polluted environment, but the actual [HNO ₃ ] would be even lower because nitric acid ionized in solution. In other words, once nitric acid is in solution, it will not come back out again unless the droplet evaporates; conversely any vapor-phase nitric acid will be quickly absorbed into the aqueous-phase in the presence of cloud or fog water. Which pollutants can be rained out?

What is the possible pH of water in a high cloud (alt. ≃ 5km) that absorbed sulfur while in equilibrium with 100 ppb of SO ₂ ? The pressure decreases as a function of height. At 5km the ambient pressure is around half the atmospheric pressure: 0.54 atm. This SO ₂ will not stay as SO ₂ H ₂ O, but participate in a aqueous phase reaction, that is it will dissociate.

The concentration of SO ₂ H ₂ O, however, remains constant because more SO ₂ is entrained as SO ₂ H ₂ O dissociates. The extent of dissociation depends on [H ⁺ ] and thus pH, but the concentration of SO ₂ H ₂ O will stay constant as long as the gaseous SO ₂ concentration stays constant. What’s the pH for our mixture? If most of the [H ⁺ ] comes from SO ₂ H ₂ O dissociation, then Note that there about 400 times as much S in the form of HOSO ₂⁻ as in the form H ₂ OSO ₂. HOSO ₂⁻ is a very weak acid, ant the reaction stops here. The pH of cloudwater in contact with 100 ppb of SO ₂ will be 4.5

Because SO ₂ participates in aqueous-phase reactions, Eq. (I) above will give the correct [H ₂ OSO ₂ ], but will underestimate the total sulfur in solution. Taken together all the forms of S in this oxidation state are called sulfur four, or S(IV). If all the S(IV) in the cloud water turns to S(VI) (sulfate) then the hydrogen ion concentration will approximately double because both protons come off H ₂ OSO ₄, in other words HSO ₄⁻ is a strong acid. This is fairly acidic, but we started with a very high concentration of SO ₂, one that is characteristic of urban air. In more rural areas of the eastern US an SO ₂ mixing ratio of a 1-5 ppb is more common. As SO ₂ H ₂ O is oxidized to H ₂ OSO ₄, more SO ₂ is drawn into the cloud water, and the acidity continue to rise. Hydrogen peroxide is the most common oxidant for forming sulfuric acid in solution.