Acidification Processes in Soil and Bedrock

An acid is a compound which has the capability to release H+ ions (Box 1). High concentrations of H+ (low pH values) can attack natural materials such as limestone, soil minerals, and living tissues or man-made materials or artwork such as cement, concrete, metal surfaces, or sculptures.

Acidification of soils is a natural process on geological timescales. In general, soil acidification can be described as a two-step process:

1. The slow gradual depletion of nutrient cations, that is, the leaching of Ca2+, Mg2+, K+, bases (HCO3, CO3~, etc.).

2. Their replacement by 'acidic' H+, Al, Fe, and Mn ions or complexes. While H+ is mainly supplied by internal ecosystem processes or by atmospheric deposition, the 'acidic' metal cations are released from the bedrock by mineral weathering.

Intensive agriculture and forestry can lead to high ecosystem internal H+ production (see below). Many man-made landscapes originate from extensive land-use

Box 1 H+ ion concentration or pH

The pH scale is logarithmic, neutral water has a pH of 7.0. In the absence of strong acid anions such as SO43 and NO3, pure rainwater has a pH of 5.6-5.7. This means that 'clean' rain in equilibrium with atmospheric carbon dioxide (CO2) is acid.

Each whole unit on the pH scale represents a multiplication factor of 10. Thus, water with a pH of 5.0 is 100 times more acidic than water with a pH of 7.0.

activities. One example is the heath lands in northwestern Europe, where human pasture, field, and forest management over centuries have led to soil acidification and erosion. Plant material was removed by grazing. Sparsely growing trees were used in salt refineries and as firewood. The humus layer including the ground vegetation was removed and interspersed in stables. The mixture of soil and dung was used for manuring fields at locations different from the areas where organic material had been removed.

Application of dung, liquid manure, or compost can compensate (part of) nutrient losses. Agriculture and partly also forestry often apply multinutrient mineral fertilizers containing lime (H+ buffering CaCO3/MgCO3).

Most crystalline shields and noncarbonated sedimentary rocks can be considered as being sensitive to acidification by 'acid rain'. Areas where acidification has been an issue are major parts of northern Europe, northeastern USA, eastern Canada, and parts of China. Due to rapid increases in acidifying emissions potential future problem areas could be Nigeria, India, Venezuela, Southern Brazil, and Southeast Asia.

Hydrogen Ion (H+) Sources to Soils

Several biologically mediated processes lead to ecosystem internal H+ production (=H+ sources), while atmospheric deposition or mineral fertilizer applications are external H+ sources.

Nutrient cation uptake and the consequences of biomass export

The majority of nutrients needed for plant growth exist as cations (Ca2+, Mg2+, K+, Na+, NH4, Fe+, etc.). Fewer nutrients and their less amounts are taken up as anions (NOf, HPO^, SO4~, etc.). The latter implies that vegetation assimilates an excess of non-N nutrient cations over anions. To compensate for electroneutrality, plants release either weak organic acids or H+ to the soil solution for each positively charged ion taken up by the roots (e.g., one H+ in the case of K+ uptake and two H+ for each Mg2+ ion). As a result the pH of the soil solution near the root surface (the rhizosphere) can drop considerably during the growing season. However, if no plant material is removed from the system, nutrient cations return to the soils during decomposition, which is an H+-consuming process. Thus, without biomass export, plant uptake has no long-term effect on acidification.

In a growing forest the consequence of nutrient cation uptake poses a net production of H+ in the soil solution over decades, because nutrients can be stored in the biomass and humus layer for a relatively long time periods (Figure 1). However, organic matter in natural ecosystems is usually exchanged in cycles, that is, when a forest or part of a forest dies, assimilated nutrients are released again via decomposition.

Humus layer or top soil: accmulation of nutrients

Humus layer or top soil: accmulation of nutrients

Mariculture Humus

Root

soil soil

Mineral soil: loss of nutrients and accumulation of H+ ions

Root

soil soil

Figure 1 Nutrient uptake. NC+, nutrient cations; A , anion.

Acid rain

Forestry Animal I Biomass

Agriculture husbandry Í* export

mmi

i

aSpppMgg

{

T

Depletion of nutrient cations to ;¿production,of .acidity -

Under natural conditions, the deprotonation of H2CO3 is the most significant H+ source in acidifying soils down to pH > 5 (note that the pKa of H2CO3 is 6.46 preventing a decrease of pH below 5). Thus, CO2 is the major agent of CO2~, mineral weathering, and natural acidification (see the section entitled 'Hydrogen ion (H+) sinks').

Figure 2 Biomass export.

Below pH 5, acidification.

production organic acids drive natural

Thus, excess uptake of positively charged nutrients by plants affects soil acidity in the long term only if plant material is removed (Figure 2). This removal can be driven by harvesting grain crops in agriculture, by removing cattle, which have converted part of the plant material they have eaten into body tissue, or by using timber and firewood in forestry.

Decomposition, root respiration, and the production of carbonic acid

Microbial degradation (decomposition) of organic material and root respiration lead finally to relatively high CO2 concentrations in the soil air (high CO2 partial pressure). A greater part of this CO2 resolves in soil water and forms carbonic acid (Box 2). A consequence is that waters percolating through soils (or bedrocks) contain usually substantially higher concentrations of H2CO3 than rainwater or surface waters. The major anion produced by H2CO3 dissociation is HCO^ (Box 2).

Soil organic matter and the production of soil organic acids

Soil organic matter consists of carbohydrates, which contain acidic groups (e.g., carboxyl, carbonyl, or hydroxyl). An increase in soil organic matter is in itself a potential source of acidity, as also the application of dung or liquid manure. However, organic matter contains only weak acids, that is, in contrast to strong acids such as H2SO4 they do not dissociate completely but release only a portion of their H+. This proportion varies according to the H+ concentration in the solution. The lower the pH, the fewer the H+ ions released (and the more the acid groups protonated).

The deprotonation of dissolved organic acids can be described by dissociation constants (pKa values). The lower the pKa, the stronger is the acid (Table 1). Dissolved organic acids are ubiquitous in soils and can deprotonize depending on the pH values.

Transport of organic anions causes soil acidification in deeper soil horizons (a process called podsolization) and

Table 1 pKa values of some important inorganic and organic acids

Acid Formula pKa1 pKa2 pKa3

Sulfuric acid H2SO4 -3 1.92

Nitric acid HNO3 -1.32

Oxalic acid (COOH)2 1.23 4.19

Phosphoric acid H3PO4 2.12 7.21 12.67

Formic acid HCOOH 3.75

Acetic acid CH3COOH 4.75

Carbonic acid H2CO3 6.46 10.25

Humic and fulvic acids Complex organic molecules in soil solution and freshwaters; pKa between 3 and 8

waters. In general, organic anions can be rapidly degraded to CO2 by microbial activity and they are important components of groundwater or stream water acidification only in fens or bogs.

H+ turnover within the nitrogen cycle

The N cycle (Figure 3) is connected to major H+ turnover processes in soils. Nitrogen is one of the major nutrients, and N turnover exceeds the turnover of all other nutrients and trace elements quantitatively, with the exception of carbon. Because inorganic N can occur as a cation or an anion in soils, the influence on H+ budgets caused by N turnover is complex (Box 3).

Decomposition of N-containing organic material is usually followed by the oxidation of NH+ (nitrification), which is connected to the production of 2 mol of H+ for each NH+ molecule (Box 3).

Nitrogen is an important nutrient as it is part of proteins and nuclides in living organisms. The assimilation of NH+ during the production of amino acids produces 1 mol of H+ per mole of NH4.

Adding N as NH+ fertilizer will cause acidity ((NH4)2SO4 or NH4NO3). If the NH+ added is converted to NO- and leached out of the soil, then a very rapid rate of acidification occurs. If plants take up the NH+, then an intermediate rate of acidification occurs. Fertilization with NH 3 can actually be a neutral process (Box 4).

Leaching Figure 3 Nitrogen cycle.

NH3-volatilization

Leaching Figure 3 Nitrogen cycle.

NH3-volatilization

However, fertilizing NH3 is bound to extreme rates of volatilization and will increase local and regional N deposition dramatically.

Another way for N into the ecosystem is the fixation of N2 from the atmosphere by bacteria. While there are only few free living species, most N-fixing bacteria live in a symbiosis with plants, for example, within a legume nodule where the very stabile molecular N2 is converted into a form available for plants to use. N2 fixation involves no H+ transfer. Only after ecosystem internal mineralization of organic N to NH4 or NO-, will fixed N2 become involved in H+ transfer. If N accumulates in the ecosystem, it usually does so in soil organic matter. Besides leaching and harvest as an N loss to ecosystems, N can volatilize into the atmosphere as gaseous N compound via denitrification (bound to consumption of H+) or volatilization (production of H+).

To conclude, a disruption of the N cycle by either biomass export (harvest) or fertilization has major consequences concerning the H+ balance and thus the acidification of soils.

Oxidation of reduced compounds

Increasing water saturation promotes anoxic conditions in soils, and microorganisms can use NO-, SO4-, Fe, Mn, and CO2 as electron acceptors instead of O2. During such reduction processes, H+ is consumed and alkalinity is generated. Up to 70% of the impacted acidity can be neutralized in forested freshwater wetlands by Fe and SO4- reduction alone.

Conversely, during the oxidation process of (prior) reduced compounds, H+ is released. For example, if soils or waters contain a substantial amount of reduced Fe and have a low buffering capacity, the pH of the solution may fall from a value of about 6-7 to 2-3 caused by the oxidation of formerly reduced Fe compounds. Concerning SO4-reduction, FeS and FeS2 are the most important products. If conditions stay anoxic over a longer time period, reduced S species might be incorporated into the organic substance and thus stored long term, resulting in an equally long-term alkalinity generation.

Box 3 Proton sources and sinks within the nitrogen cycle (blue, H+ sink; red, H+ source)

N2-fixation:

Ammonification:

N2 ! Norg No ions in solution, thus no H+ turnover

[1]

Denitrification:

R-C-NH2 + H2O + H+ ! NH+ + R-C-OH

[2]

NO- uptake:

2NO3 + 3CH2O + 1 /2O2 + 2H+ ! N2 + 3CO2 + 4H2O

[3]

NH|-uptake and assimilation:

NO3- + H+ ! Norg

[4]

Nitrification:

NH+ + R-C-OH ! R-C-NH2 + H2O + H+

[5]

NH3-volatilization:

NH+ + 2O2 ! H2O + NO3- + 2H+

[6]

aNote that consumption of OH3

NH+ + OH-a ! NH3 + H2O is equivalent to production of H+ and vice versa.

[7]

Box 4 Acid production due to nitrogen

fertilization (negative for H consumption, positive

for H+ production)

Application of ammonium (e,g., as(NH4)2SO4)

NH4 uptake and assimilation:

+1 mol H+ per NH+

Nitrification:

+2 H+ mol H+ per

NH+

If produced NO3 is taken up by plant:

-1 mol H+ per NO-

Application of NH4NO3

NH4 uptake and assimilation:

+1 mol H+ per NH+

Nitrification:

+2 H+ mol H+ per

NH+

If produced or applied NO3 is taken

-1 mol H+ per NO-

up by plant:

Application of Ammonia (NH3)

Dissolution (re-volatilization):

-1 mol H+ per NH3

Nitrification of produced NH4:

+2 H+ mol H+ per

NH+

If produced NO3 is taken up by plant:

-1 mol H+ per NO-

Drainage of valley floors and thus exposure to air (O3) causes reduced compounds to re-oxidize and release substantial amounts of acid. In contrast, wetland soils and riparian zones may act as long-term sinks for deposited H+, SO2~, and NO 3, depending on soil characteristics, climatic parameters, and the composition of the soil microbiota.

Atmospheric deposition of acidifying compounds

Acidifying pollutants are deposited into ecosystems as follows:

1. directly as gases and aerosols to vegetation or other surfaces (dry deposition, especially NHj);

2. as rain or snow (wet deposition); and

3. via impaction and sedimentation of fog or droplets to various surfaces (occult deposition).

High acidification rates occur in forested coniferous sites (compared to deciduous sites) due to more efficient scavenging of acidifying pollutants from the atmosphere especially during wintertime. The acidification rate caused by acid deposition is in the range 0.8-7 kmol ha-yr-1 due to the combined effects of HNO3, H2SO4, HCl, and NH4 deposition.

Hydrogen Ion (H+) Sinks

In natural ecosystems weathering of minerals counteracts acidification, that is, acts as H+ sink. Thus, main sinks in ecosystems are geochemical buffering reactions in soils and bedrocks, and only a minor fraction is buffered in waters. The so-called acid neutralization capacity in soils and bedrock (ANCsolid) can be defined as the sum of all unprotonated buffering substances. Thus, acidification is always accompanied by a decrease in ANCsolid over time. It is important to note that this decrease in ANCsolid is irreversible (against the background of our human calculation of times).

Besides the capacity, that is, the total buffering pool of a soil or bedrock, the geochemical reaction rate of the buffering substances are a crucial factor determining how much of the acidifying compounds are neutralized over a certain period. This rate can, for example, be estimated as kilomole charge per hectare per year (kmolc ha— yr— ).

Ecologically effective are, last but not the least, the concentrations of certain ions in the soil/bedrock solution or freshwaters. Such intensity parameters can be measured as concentrations (e.g., pH, Mg2+, or Al3+; in molc per liter) or as base cation saturation of exchanger complexes in soils (in %).

In theory, many of the geochemical buffering processes are equilibrium reactions. However, the loss of ions with percolating water leads to permanent disequilibria. The latter has important implications for the expected reversibility of soil acidification under decreasing deposition regime. In North America and Europe, soil acidification is irreversible as long as the supply of weathering products from bedrock is smaller than the loss of weathering products due to the combined effects of natural and anthropogenic acidification processes. It is highly unlikely that the latter two will become smaller than the supply by weathering products because natural acidification processes already exceed buffering by weathering processes in most systems.

Carbonate dissolution

The H+ ions are buffered via the dissolution of Ca(or Mg)CO3 in soils and bedrock (Box 5) as long as soils or bedrocks contain accessible carbonate.

The pH values in the soil solution are quasiconstant and stay above pH 6.2 (Figure 4). The buffering rate is high, for example, 2 kmolc ha— yr— at a water percolation rate of 2001 m— and a CO2 partial pressure of 0.3 kPa CO2. The CO3— buffering is usually an irreversible oneway reaction resulting in the loss of Ca2+ and HCO— from soils and bedrock.

Buffering substance/process: Carbonates

Silicates pH

Variable loading

Oxides/hydroxides

1 1 Range where process is regulating solution pH Figure 4 Buffering systems.

Box 6 Weathering of primary silicate minerals

—(SiO)M + H+ )—(SiOH) + M+ (M+ = exchangeable metal cation)

Compared to the accumulated H+ production or input rates in ecosystems affected by acid deposition (see the section titled 'Hydrogen ion (H+) sources in soils') this rate is rather low.

Silicate weathering results usually in the irreversible destruction of clay minerals, the release of exchangeable cations, and Al ions to the soil solution. Dominant anions in the solution are HCO— and organic anions.

Further, some clay mineral crystals can dissolve completely, leading to high Al concentrations in the solution (Box 7).

Silicate weathering

The H+ ions are buffered by the (relatively slow) weathering of primary silicate minerals (e.g., Box 6). The soil solution stays in the pH range of 6.2-5.0 (Figure 4), and the rate lies between 0.2 and 2 kmolc ha— yr— .

Box 5 Dissolution of calcium carbonate

CaCO3 + H+ ) HCO3- + Ca+ pKai = 6.46 HCO3- + H+ ) CO2 + H2O pKa2 = 10.25

Exchanger with variable loading

The H+ ions exchange against base cations bound to clay minerals and oxides or organic matter (pH range 5-4.2; Figure 4; Box 8).

Exchangeable cations are lost with the percolating water. The buffering capacity depends on the absolute cation exchange capacity and on the percentage ofsatura-tion of the exchanger complex with base cations. The buffering rate is very high (fast reaction). The Ca2+ ion is usually the dominant cation. In systems influenced by 'acid rain', HCO3 is replaced by the anions SO2+ and NO —. In naturally acidic ecosystems, for example, in bogs, organic anions are dominant.

Box 8 Exchange of H+ ions against base cations (clay minerals, oxides, organic matter)

(M+ = exchangeable metal cation)

(pKa dependent; pKa 3-8)

Amorphous hydroxides or oxides of aluminum and iron

The so-called Al and Fe buffer ranges can be described by the equilibrium reactions shown in Box 9.

The H+ ions are bound in water (H2O), and soluble Al (or Fe) ions (ion complexes) emerge in the soil solution. The neutralization capacity depends on the reactive amount of Al or Fe (hydr)oxides. Buffering rates are high, and Al ions (or Fe ions) become the predominant cations in the soil solution. In ecosystems affected by acid deposition, SO4- and NO- are the predominant anions.

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