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BUFFERING CAPACITY
142 BUFFERING CAPACITY
BUFFERING CAPACITY B R James, University of Maryland, College Park, MD, USA
The Samovar Analogy for Buffering in Soils
ß 2005, Elsevier Ltd. All Rights Reserved.
The concepts of the intensive variable, capacity factor, buffer capacity, and buffering mechanisms for soils can be modeled and visualized by analogy with old-fashioned, Russian samovars used to make large volumes of tea (Figure 1). Each samovar comprises a large copper reservoir to boil water and a narrow glass tube on the outside of the samovar’s tank to indicate how full the tank is. As boiling water is drawn out of the tank through the spigot, the indicator tube empties quickly and temporarily, but is refilled to the new tank level when the spigot is shut off (as indicated by the arrows between the tank and indicator tube in Figure 1). To decrease the volume of a samovar tank by a given fraction of its total capacity (e.g., from level 1 to level 2 in the indicator tube), a greater volume of water must be withdrawn from the larger samovar than from the smaller one. Since the level of water in the indicator tube is proportional to how full each tank is, it is a parameter that is independent of the volume of the samovar. The height of the water column is analogous to an intensive variable in soil solution such as pH, a measure of how acidic an aqueous sample is, but not an indicator of its volume. The volume of the samovar represents the capacity factor (e.g., total soil acidity), and the
Introduction Buffering is the resistance of a system to change in response to a perturbation, and it is a key attribute of soils from the molecular level to the landscape scale. Chemical, physical, and biological processes may raise or lower solute concentrations (or activities) as intensive variables in soil solution, thereby temporarily disturbing dynamic equilibria or steadystate conditions between soil water and solid or gas phases of the soil. In response to such a perturbation, one or more processes may release solute to soil solution or remove it to restore wholly or partially the original concentration. The reservoir of such solutes in solution, in solid phases, or in the soil atmosphere for restoration of soil solution chemical composition is known as the ‘capacity factor.’ The measurable change in the capacity factor per unit change in the intensive variable is called the ‘buffer index,’ ‘buffer intensity,’ or, more commonly, ‘buffer capacity.’ Myriad chemical reactions (buffering mechanisms) govern such release or uptake of ions and molecules between soil solution and solid or gas phases, including cation and anion exchange, oxidation–reduction reactions, dissolution–precipitation processes, and metal–organic ligand complexation. Understanding and quantifying buffer capacities and buffering mechanisms at the colloid and molecular level of soils can aid in predicting the sensitivity and resilience of soil-water systems to anthropogenic and natural perturbations of ecosystems. New ecological theory related to succession dynamics of disturbed ecosystems identifies that natural systems, including soils, are constantly and naturally recovering from regular and irregular disturbances of various severities. The processes and mechanisms of recovery determine the biodiversity and stability of ecosystems, and buffering in soils and natural waters is a key control of nutrient availability and pollutant bioavailability. Buffering reactions influence water quality and community regrowth, migration, and recruitment in disturbed patches on the landscape. The success of human efforts to restore disturbed ecosystems and soils is also determined in part by the scale and nature of soil-buffering processes, as related to natural ones that govern element transformations.
Figure 1 Samovar models representing small and large buffer capacities (tank sizes) in equilibrium with the same level of intensity variable (represented by the height of liquid in the narrow indicator tube outside the tank). (Adapted from Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall.)
BUFFERING CAPACITY
Figure 2 Samovar models representing small and large buffer capacities (tank sizes) in equilibrium with the same level of intensity variable (represented by the height of liquid in the narrow indicator tube outside the tank). The larger tank with the hourglass shape models a buffered system in which the buffer capacity (change in tank volume per unit change in the height of the liquid in the indicator tube) is not constant and varies during titration (drainage of the tank). (Adapted from Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall.)
change in the tank volume per unit change in the level of the indicator tube is the buffer capacity. The restoration of the water level in the indicator tube from the samovar after shutting off the spigot is due simply to the flow of water, i.e., the buffering mechanism in this model. In this simple comparison of two samovars of different volumes, but with identical fractions of their capacities filled with water, as the boiling water is drained completely, the buffer capacity (change in tank volume per change in water level) remains constant. In complex soil colloid-water systems, however, buffer capacities do not usually remain constant as an intensive variable is changed, and the magnitude of the buffer capacity will change accordingly. This is modeled by a peculiar samovar with a narrow middle section and wide sections at the top and bottom (Figure 2). In this samovar model, the system would be well buffered at the beginning and end of draining the tank, but would be poorly buffered in the middle. In soils, measurements of intensive variables are often made routinely (e.g., pH, Ca concentration, partial pressure of O2, and oxidation–reduction potential), and are used to indicate the energy state or ‘ability of the system to do work’ in affecting other coupled systems such as groundwater or plants rooted in the soil. The size of the capacity factor and buffering mechanisms linking the intensive and capacity parameters
143
are often not known with the same accuracy as are measurements of intensive variables in soils (i.e., the samovar tank is invisible). Therefore, investigations of the nature and scale of buffering processes in soils are needed to understand controls on the value of an intensive variable and what chemical and biological processes control its resistance to change. The samovar analogy for soil buffering will be used to explain similarities and differences among several soil chemical processes responsible for buffering of soil acidity, oxidation–reduction status, and metal ion activities controlled by dissolution–precipitation and organic complexation by ligands. Soil acidity and oxidation–reduction status of soils are considered master variables that control the steady-state chemical composition of soils and the kinetics of approach to chemical equilibrium. As a result, predicting their resistance to change in response to disturbance of soil conditions is relevant to many environmental quality, plant nutrition, and human health impacts of soils. Dissolution and precipitation reactions are important in weathering, soil development, and leaching processes, as well as in controlling nutrient and pollutant solubility, speciation, bioavailability, and cellular absorption. Organic and inorganic ligand complexation, typically of cations, controls ion activity in equilibrium with soluble and insoluble forms of the element. Each of these processes is a coupling of an ion activity (intensive variable), a capacity factor (reserve of ion), and a buffering capacity and buffering mechanism that govern how the intensity and capacity factors interrelate in a soil system.
Buffering of Soil Acidity The intensive variable for soil acidity is conceptually defined as the hydrogen ion activity of soil pore water or ‘active acidity,’ and is operationally defined by measuring pH in a soil-water or dilute salt suspension with a glass-reference electrode system, where pH is log (Hþ) and (Hþ) is the hydrogen ion or proton activity in units of moles per liter. To neutralize the active acidity (e.g., raising pH from 5.0 to 7.0) in the soil solution of the top 15 cm of a hectare of loamy soil at field capacity moisture (10 kPa water potential) containing 200 g montmorillonite and 30 g organic matter per kilogram of soil would require approximately 500 g of CaCO3. This is just a handful of lime, compared with the 10 000 kg CaCO3 ha1 that might be needed to raise the whole soil pH to 7.0. The 20 000:1 ratio of lime needed to neutralize the capacity factor or ‘reserve acidity’ of the soil to that needed to neutralize the active acidity reflects the relative size of the metaphorical samovar tank and its indicator tube for soil acidity.
144 BUFFERING CAPACITY
Soil properties that determine the buffer capacity for soil pH include initial pH, aluminosilicate clay content, and mineralogy; organic-matter content; and Al(III), Fe(III), and Mn(III,IV)(hydr)oxide contents. Soil pH buffer capacity is determined by adding increasing quantities of acid or base to a given mass of soil and measuring pH after allowing sufficient equilibration time for pH to stabilize. It is often designated: ¼ dC=dpH
½1
where is buffer index or capacity in units of moles of Hþ or OH added per kilogram of soil (dC) per unit pH change, dpH. The buffer capacity is quantified as the reciprocal of the slope of a linear relationship between measured pH plotted on the ordinate versus moles of Hþ or OH added per kilogram of soil on the abscissa. Other, quicker indicator tests for reserve or total acidity as the capacity factor have been developed to estimate ‘lime requirement’ for acidic agricultural soils used to grow crops that are sensitive to acid soil conditions. Examples of these ways to estimate reserve acidity include chemical extractions of exchangeable and reactive Al3þ, measures of cation exchange capacity (CEC) and the fraction of exchange sites occupied by ‘basic’ cations (predominantly Ca2þ, Mg2þ, Kþ, and Naþ), measurements of decreases in the pH of well-buffered solutions added to the soil sample, and equilibrations of a soil sample with a BaCl2 solution buffered at pH 8.2, followed by titration to pH 5 to determine the quantity of acid that reacted with the buffer. Measures of free CaCO3 content can be used to estimate the total alkalinity of a soil and its buffer capacity against acidification. Diverse chemical reactions occurring on colloid surfaces of clay, oxides, and organic matter are responsible for pH buffering. They involve cation exchange reactions, dissolution and precipitation of sparingly soluble compounds, and surface charge changes in response to pH. In most soils, pH buffer capacity and buffering mechanisms are due to combinations of these reactions, and titrations and quick tests for total acidity do not distinguish between them. None the less, understanding the relative importance of these processes and their chemical basis allows predictions of the buffering behavior of diverse soils, and it allows extrapolation to unstudied soil systems and environments, particularly as related to ecological processes, soil contamination, and soil remediation. Dissolution and Precipitation of CaCO3 and Al(OH)3
Soils containing free CaCO3 have a pH between 7.0 and 9.5 (typically 8.0–8.5) as governed by the
dissolution and precipitation of CaCO3 in water containing partial pressures of CO2 as high as 10 times that of atmospheric levels. Acidity produced by the hydration of dissolved CO2 and by other acid-generating reactions reacts with CaCO3 in accordance with: CaCO3 þ H3 Oþ ¼ Ca2þ þ HCO 3 þ H2 O
½2
In the presence of high levels of soluble or exchangeable Mg2þ or Naþ, soluble MgCO3 or Na2CO3 will form and allow the soil pH to rise to values higher than that controlled by CaCO3–CO2 equilibria (sometimes to pH >10). Addition of base or acid will be buffered by the lime in the soil, thereby maintaining its pH at approximately 8.2. Under strongly acid soil conditions, Al3þ in soil solution and on cation exchange sites hydrolyzes and generates acidity, while the dissolution of Al(OH)3 neutralizes it, leading to large values at pHs <4.5: Al3þ þ 3H2 O ¼ AlðOHÞ3 þ 3Hþ
½3
Soluble and exchangeable Al3þ is released from aluminosilicate clays upon weathering of Si-dominated tetrahedral sheets (releasing soluble monosilicic acid, Si(OH)4), thereby exposing Al-dominated octahedral sheets to Hþ attack. As a result, the buffer capacity of strongly acid soils (pH <5) is dominated by Al chemistry. Cation Exchange Reactions and pH Buffering
In the intermediate pH range between 4.5 and 6.5, cation exchange reactions on permanent aluminosilicate cation exchange sites, and on pH-dependent sites of organic matter and (hydr)oxides govern uptake and release of Hþ and OH as controls on the buffering-capacity size and mechanism for pH control. Permanent, negatively charged cation exchange sites on aluminosilicate clay minerals are due to isomorphous substitution of lower charge cations for Si4þ (e.g., Al3þ) in tetrahedral sheets and for Al3þ (e.g., Mg2þ and Fe2þ) in octahedral positions. Cation exchange sites created in this way during clay mineral formation are dominated by Ca2þ at and above pH 7 and become increasingly satisfied by Hþ and Al3þ at lower pH values. In this way, the so-called base saturation (an intensive variable reflecting the fullness of the acidity samovar) decreases as exchange sites are occupied by Al3þ and Hþ. Organic matter functional groups (principally carboxylic acids, represented by R COOH, and phenw olic acid groups, such as OH on aromatic rings) w are weak acids with association constants (pKas) in the range of 2–7 for carboxylic acids and 7–10
BUFFERING CAPACITY
for phenolic acid groups. These groups deprotonate at pHs >pKa, and protonate at pHs
Oxide and hydroxide coatings of Fe(III), Mn(III,IV), and Al on sand, silt, and clay minerals are dominated by oxygen- and hydroxide-rich planes exposed to soil solution, and changes in pH protonate and
145
deprotonate these functional groups, thereby creating variable or pH-dependent charge colloid surfaces that contribute to buffer capacity: 0 III MeIII OHþ1 2 þ OH ¼ Me wOH þ OH w ¼ MeIII O1 w
½4
where Me(III) represents a trivalent metal structurally associated with the mineral surface, and the charges on the surface become increasingly negative as pH is raised. Highly weathered soils that contain few aluminosilicate clay minerals or organic matter are often dominated by these variable-charge colloids, and they contribute to pH buffering through positive- and negative-charge creation. The relative importance of CaCO3 dissolution and precipitation, Al hydrolysis, cation exchange on permanently charged aluminosilicate minerals, organic-matter functional group protonation and deprotonation, and variation in negative and positive charges on variably charged oxide coatings depend on the relative abundance of these colloids in a given soil. In many soils, more than one of these soil chemical processes are responsible for pH buffering, and titrations of whole-soil materials reflect an integration of the several processes. Table 1 summarizes the relative magnitudes of the buffer capacities associated with these processes.
Table 1 Approximate buffer capacities of soil materials and horizons in the pH range 3.0–10.0
Soil constituent or horizon
Buffering pH range
Buffering index (mmol kg 1 pH unit 1)
Aluminosilicate clays Smectite Vermiculite Illite Kaolinite Soil organic matter
3–10 3–10 3–10 3–10 3–10 5–8
178–333 333–444 44–88 2–11 360–444
Allophane and imogolite
3–10
44–111
Fe(III) and Al(III) (hydr)oxides CaCO3 and MgCO3 Forest floor organic horizons Mineral horizons of agricultural and forest soils
3–10
11–89
>7 3–7
4444 180–360
4–10
53
pH-buffering mechanism
Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Protonation–deprotonation of weak acid functional groups; conformational changes Protonation–deprotonation of weak acid functional groups; conformational changes Protonation–deprotonation of weak acid functional groups; conformational changes Dissolution/precipitation Metal–ligand complexation/decomplexation; cation exchange; protonation/deprotonation; of weak acid functional groups Cation exchange, mineral dissolution/precipitation, Al hydrolysis, protonation–deprotonation of weak acid functional groups, conformational changes, and metal–ligand complexation/ decomplexation; principally associated with soil organic matter
Source: Bloom PR (2000) Soil pH and pH buffering. In: Sumner M (ed.) Handbook of Soil Science, pp. 333–352. Boca Raton, FL: CRC Press, with permission; McBride MB (1994) Environmental Chemistry of Soils. New York, Oxford University Press, with permission; James BR and Riha SJ (1986) pH buffering in forest soil organic horizons: relevance to acid precipitation. Journal of Environmental Quality 15: 229–234, with permission.
146 BUFFERING CAPACITY
Buffering of Soil Oxidation–Reduction Status The concept of ‘electron activity’ of soils refers to the thermodynamic tendency for electrons to be transferred from reductants (e.g., Fe(II)) to oxidants (e.g., O2), and it is operationally defined by Pt electrode potentials versus a standard reference electrode. When the measured potential is corrected for the reference electrode potential relative to the hydrogen electrode (E0 ¼ 0.0 V), it is designated Eh. This voltage can be converted to pE, a parameter analogous to pH, by dividing Eh in volts by 0.059 (the Nernstian slope factor relating electrode voltage to electron activity). The Eh or pE is the intensity variable that is a measure of electron activity or ‘electron pressure’ in a soil system. Buffering of electron activity is called poise, and is analogous to proton buffering, except that the processes responsible for resistance to change in pE are due to electron transfer reactions, many of which are microbial and enzymatically catalyzed. Heterotrophic microbial respiration in soils uses organic C as the electron source (reductant) and an array of electron acceptors as oxidants. Strict aerobes use dissolved O2 as the required oxidant to derive energy from the oxidation of organic C, the most energy-efficient process for cellular respiration. In flooded soils in which all pores are filled with water, the diffusion rate for dissolved O2 is 10 000 times slower than it is in air. As a result, microbial respiration may deplete available O2 faster than it is replenished, thereby leading to the onset of anaerobic conditions, with a decrease in the measured pE. Under anaerobic conditions, facultative anaerobes and strict anaerobes use alternative electron acceptors, with decreasing metabolic efficiency in the order (depending on pH): NO 3 , Mn(III,IV)(hydr)oxides, Fe(III)(hydr)oxides, SO2 , CO2, and Hþ. 4 The quantities of each of these electron acceptors (the capacity factor for electron activity) in the soil will ‘poise’ the Eh at a given level as governed by the free energy of reaction associated with its reduction, as coupled to the oxidation of organic C. The expected pE and Eh values at which each of these couples would poise the soil are shown at pH 5 and 7 in Table 2. Since the molecular reaction mechanism for many reduction reactions involves the transfer of the electron with a proton to the oxidant (equivalent to the addition of a H atom), the overall reaction raises the soil pH in many cases. As a result, the higher the pH of a soil, the lower the Eh or pE at which a given reduction reaction is expected to occur and at which the soil system will be poised (i.e., at a lower (Hþ), a greater (e), or ‘electron pressure’ is needed to effect the given reduction).
Table 2 Common reduction half-reactions that poise soils at designated pE values at pH 5 and 7; the log K value is the pE for the reaction at pH 0 Poising pE Half-reaction (for 1 electron reduction)
þ
1/4O2 þ e þ H ¼ 1/2H2O þ 1/5NO 3 þ e þ 6/5H ¼ 1/10N2 þ 3/5H2O 1/2Mn3O4 þ e þ 4Hþ ¼ 3/2Mn2þ þ 2H2O Fe(OH)3 þ e þ 3Hþ ¼ Fe2þ þ 3H2O þ 1/8SO2 4 þ e þ 5/4H ¼ 1/8H2S þ 1/2H2O 1/8CO2 þ e þ Hþ ¼ 1/8CH4 þ 1/4H2O e þ Hþ ¼ 1/2H2 1/4CO2 þ e þ Hþ ¼ 1/24C6H12O6 þ 1/4H2O
log K
pH 5
pH 7
20.8 21.1
15.6 14.3
13.6 11.9
30.7
16.7
8.7
15.8 5.2
4.8 1.0
1.2 3.5
2.9 0 0.21
2.1 5 5.9
4.1 7 7.9
Buffering of Ion Activities via Dissolution–Precipitation, Ion Exchange, and Ligand Complexation The soil solution activities of ions other than Hþ and OH are intensity measurements that may be controlled by dissolution–precipitation, ion exchange, and ligand complexation reactions responsible for buffering in soils. Solubility, exchange, and complexation equilibria maintain ion activities in soil solutions as leaching, and plant and microbial uptake; and other biotic and abiotic processes occur and increase or decrease particular ion activities. The disturbance of these equilibria will induce restorative shifts in accordance with the LeChatelier principle, which states that the balance of products and reactants in a dynamic chemical equilibrium will shift in response to a perturbation in order to restore the original balance of reactants and products. For example, if sparingly soluble PbCrO4 is present in a soil, Pb2þ and CrO2 concentrations of approxi4 mately 5.3 107 mol L1 will be maintained in soil solution as controlled by the solubility product, Ksp, of 2.8 1013. Any process that depletes the soluble Pb2þ or CrO2 4 will induce the dissolution of a small amount of solid PbCrO4 to restore the equilibrium activities of the Pb2þ and CrO2 4 . Similarly, additions of these ions to the soil solution will induce precipitation of more PbCrO4. In this way, the ion activities are buffered and maintained in soil solution, as shown in eqn [5]: PbCrO4 ¼ Pb2þ þ CrO2 4
½5
Similar reactions occur via cation and anion exchange on charged colloids to buffer ion activities,
BUFFERING CAPACITY
as discussed above for soil acidity buffering. The suite of exchangeable cations in soils is dominated by Ca2þ, Mg2þ, Kþ, Na2þ, Hþ, and Al3þ, depending on the pH of the soil, the mineralogy of the soil parent material, the aluminosilicate clay mineralogy, plant uptake, microbial processes, and human-induced changes in the soil conditions. Multiple equilibria are established in any soil, but all comprise a capacity factor, represented by the exchangeable cations, and an intensity factor. The capacity factor is sometimes called the ‘quantity’ term, and the quantity-to-intensity ratio (Q/I) is a measure of the nature of the cation or anion exchange buffering reactions responsible for maintaining ion activities in soil solution. Eqn [6] represents a Ca–K cation exchange reaction for such a Q/I relationship: Ex-Ca þ 2Kþ ðaqÞ ¼ Ex-K2 þ Ca2þ ðaqÞ
½6
in which Ex-Ca represents solid-phase, exchangeable Ca, Kþ(aq) is soluble Kþ, Ex-K2 is exchangeable Kþ, and Ca2þ(aq) is soluble Ca2þ. Any environmental conditions that change the activity ratio of Kþ to Ca2þ in soil solution (e.g., changing water content, preferential plant uptake of one cation over the other, precipitation of Ca2þ, or addition of Ca- or Kcontaining materials to the soil) will result in K-forCa or Ca-for-K exchange reactions that restore and buffer the ion activities in soil solution. The balance between the ratios of exchangeable Ca-to-K and soluble Ca-to-K in solution can be described by an equilibrium constant, a measure of the buffer capacity between exchangeable and soluble forms of the cations. The activities of many ions in soil solution are also in equilibrium with complexed or chelated forms of the metals that are also soluble, but in forms in which the positive charge of the cation has been neutralized by negatively charged ligands. Any processes that increase or decrease the activity of the hexaquo form of the ‘free’ form of the cation, e.g., Fe(H2O)3þ 6 , that is in equilibrium with complexed forms, e.g.,
Bulk Density
147
complexed with carboxylic acid groups, hydroxyl ions, or other Lewis bases, will induce a shift in the equilibrium to restore the balance of the complexed and free form of the ion. In this way, the ‘free’ form of the cation (its activity or intensity) is maintained through the equilibrium with the capacity factor, the complexed form. Similar to solubility products and cation exchange equilibria, a stability constant quantifies the relative thermodynamic stability of the complexed and free forms of a metal in a quantity–intensity relationship. See also: Acid Rain and Soil Acidification; Biodiversity; Environmental Monitoring; Eutrophication; Forest Soils; Microbial Processes: Environmental Factors; Nuclear Waste Disposal; Organic Soils; Pollutants: Persistent Organic (POPs); Pollution: Groundwater; Industrial; Remediation of Polluted Soils; Waste Disposal on Land: Liquid; Municipal; Wetlands, Naturally Occurring
Further Reading Bloom PR (2000) Soil pH and pH buffering. In: Sumner M (ed.) Handbook of Soil Science, pp. 333–352. Boca Raton, FL: CRC Press. Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall. James BR and Bartlett RJ (2002) Redox phenomena. In: Sumner M (ed.) Handbook of Soil Science, pp. 169–194. Boca Raton, FL: CRC Press. James BR and Riha SJ (1986) pH buffering in forest soil organic horizons: relevance to acid precipitation. Journal of Environmental Quality 15: 229–234. Lindsay WL (1979) Chemical Equilibria in Soils. New York: Wiley-Interscience. Marschner H (1995) Mineral Nutrition of Plants, 2nd edn. London, UK: Academic Press. McBride MB (1994) Environmental Chemistry of Soils. New York: Oxford University Press. Reice SR (1994) Nonequilibrium determinants of biological community structure. American Scientist 82: 424–435. Stumm W and Morgan JJ (1996) Aquatic Chemistry, 3rd edn. New York: Wiley-Interscience.
See Porosity and Pore-Size Distribution
BUFFERING CAPACITY B R James, University of Maryland, College Park, MD, USA
The Samovar Analogy for Buffering in Soils
ß 2005, Elsevier Ltd. All Rights Reserved.
The concepts of the intensive variable, capacity factor, buffer capacity, and buffering mechanisms for soils can be modeled and visualized by analogy with old-fashioned, Russian samovars used to make large volumes of tea (Figure 1). Each samovar comprises a large copper reservoir to boil water and a narrow glass tube on the outside of the samovar’s tank to indicate how full the tank is. As boiling water is drawn out of the tank through the spigot, the indicator tube empties quickly and temporarily, but is refilled to the new tank level when the spigot is shut off (as indicated by the arrows between the tank and indicator tube in Figure 1). To decrease the volume of a samovar tank by a given fraction of its total capacity (e.g., from level 1 to level 2 in the indicator tube), a greater volume of water must be withdrawn from the larger samovar than from the smaller one. Since the level of water in the indicator tube is proportional to how full each tank is, it is a parameter that is independent of the volume of the samovar. The height of the water column is analogous to an intensive variable in soil solution such as pH, a measure of how acidic an aqueous sample is, but not an indicator of its volume. The volume of the samovar represents the capacity factor (e.g., total soil acidity), and the
Introduction Buffering is the resistance of a system to change in response to a perturbation, and it is a key attribute of soils from the molecular level to the landscape scale. Chemical, physical, and biological processes may raise or lower solute concentrations (or activities) as intensive variables in soil solution, thereby temporarily disturbing dynamic equilibria or steadystate conditions between soil water and solid or gas phases of the soil. In response to such a perturbation, one or more processes may release solute to soil solution or remove it to restore wholly or partially the original concentration. The reservoir of such solutes in solution, in solid phases, or in the soil atmosphere for restoration of soil solution chemical composition is known as the ‘capacity factor.’ The measurable change in the capacity factor per unit change in the intensive variable is called the ‘buffer index,’ ‘buffer intensity,’ or, more commonly, ‘buffer capacity.’ Myriad chemical reactions (buffering mechanisms) govern such release or uptake of ions and molecules between soil solution and solid or gas phases, including cation and anion exchange, oxidation–reduction reactions, dissolution–precipitation processes, and metal–organic ligand complexation. Understanding and quantifying buffer capacities and buffering mechanisms at the colloid and molecular level of soils can aid in predicting the sensitivity and resilience of soil-water systems to anthropogenic and natural perturbations of ecosystems. New ecological theory related to succession dynamics of disturbed ecosystems identifies that natural systems, including soils, are constantly and naturally recovering from regular and irregular disturbances of various severities. The processes and mechanisms of recovery determine the biodiversity and stability of ecosystems, and buffering in soils and natural waters is a key control of nutrient availability and pollutant bioavailability. Buffering reactions influence water quality and community regrowth, migration, and recruitment in disturbed patches on the landscape. The success of human efforts to restore disturbed ecosystems and soils is also determined in part by the scale and nature of soil-buffering processes, as related to natural ones that govern element transformations.
Figure 1 Samovar models representing small and large buffer capacities (tank sizes) in equilibrium with the same level of intensity variable (represented by the height of liquid in the narrow indicator tube outside the tank). (Adapted from Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall.)
BUFFERING CAPACITY
Figure 2 Samovar models representing small and large buffer capacities (tank sizes) in equilibrium with the same level of intensity variable (represented by the height of liquid in the narrow indicator tube outside the tank). The larger tank with the hourglass shape models a buffered system in which the buffer capacity (change in tank volume per unit change in the height of the liquid in the indicator tube) is not constant and varies during titration (drainage of the tank). (Adapted from Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall.)
change in the tank volume per unit change in the level of the indicator tube is the buffer capacity. The restoration of the water level in the indicator tube from the samovar after shutting off the spigot is due simply to the flow of water, i.e., the buffering mechanism in this model. In this simple comparison of two samovars of different volumes, but with identical fractions of their capacities filled with water, as the boiling water is drained completely, the buffer capacity (change in tank volume per change in water level) remains constant. In complex soil colloid-water systems, however, buffer capacities do not usually remain constant as an intensive variable is changed, and the magnitude of the buffer capacity will change accordingly. This is modeled by a peculiar samovar with a narrow middle section and wide sections at the top and bottom (Figure 2). In this samovar model, the system would be well buffered at the beginning and end of draining the tank, but would be poorly buffered in the middle. In soils, measurements of intensive variables are often made routinely (e.g., pH, Ca concentration, partial pressure of O2, and oxidation–reduction potential), and are used to indicate the energy state or ‘ability of the system to do work’ in affecting other coupled systems such as groundwater or plants rooted in the soil. The size of the capacity factor and buffering mechanisms linking the intensive and capacity parameters
143
are often not known with the same accuracy as are measurements of intensive variables in soils (i.e., the samovar tank is invisible). Therefore, investigations of the nature and scale of buffering processes in soils are needed to understand controls on the value of an intensive variable and what chemical and biological processes control its resistance to change. The samovar analogy for soil buffering will be used to explain similarities and differences among several soil chemical processes responsible for buffering of soil acidity, oxidation–reduction status, and metal ion activities controlled by dissolution–precipitation and organic complexation by ligands. Soil acidity and oxidation–reduction status of soils are considered master variables that control the steady-state chemical composition of soils and the kinetics of approach to chemical equilibrium. As a result, predicting their resistance to change in response to disturbance of soil conditions is relevant to many environmental quality, plant nutrition, and human health impacts of soils. Dissolution and precipitation reactions are important in weathering, soil development, and leaching processes, as well as in controlling nutrient and pollutant solubility, speciation, bioavailability, and cellular absorption. Organic and inorganic ligand complexation, typically of cations, controls ion activity in equilibrium with soluble and insoluble forms of the element. Each of these processes is a coupling of an ion activity (intensive variable), a capacity factor (reserve of ion), and a buffering capacity and buffering mechanism that govern how the intensity and capacity factors interrelate in a soil system.
Buffering of Soil Acidity The intensive variable for soil acidity is conceptually defined as the hydrogen ion activity of soil pore water or ‘active acidity,’ and is operationally defined by measuring pH in a soil-water or dilute salt suspension with a glass-reference electrode system, where pH is log (Hþ) and (Hþ) is the hydrogen ion or proton activity in units of moles per liter. To neutralize the active acidity (e.g., raising pH from 5.0 to 7.0) in the soil solution of the top 15 cm of a hectare of loamy soil at field capacity moisture (10 kPa water potential) containing 200 g montmorillonite and 30 g organic matter per kilogram of soil would require approximately 500 g of CaCO3. This is just a handful of lime, compared with the 10 000 kg CaCO3 ha1 that might be needed to raise the whole soil pH to 7.0. The 20 000:1 ratio of lime needed to neutralize the capacity factor or ‘reserve acidity’ of the soil to that needed to neutralize the active acidity reflects the relative size of the metaphorical samovar tank and its indicator tube for soil acidity.
144 BUFFERING CAPACITY
Soil properties that determine the buffer capacity for soil pH include initial pH, aluminosilicate clay content, and mineralogy; organic-matter content; and Al(III), Fe(III), and Mn(III,IV)(hydr)oxide contents. Soil pH buffer capacity is determined by adding increasing quantities of acid or base to a given mass of soil and measuring pH after allowing sufficient equilibration time for pH to stabilize. It is often designated: ¼ dC=dpH
½1
where is buffer index or capacity in units of moles of Hþ or OH added per kilogram of soil (dC) per unit pH change, dpH. The buffer capacity is quantified as the reciprocal of the slope of a linear relationship between measured pH plotted on the ordinate versus moles of Hþ or OH added per kilogram of soil on the abscissa. Other, quicker indicator tests for reserve or total acidity as the capacity factor have been developed to estimate ‘lime requirement’ for acidic agricultural soils used to grow crops that are sensitive to acid soil conditions. Examples of these ways to estimate reserve acidity include chemical extractions of exchangeable and reactive Al3þ, measures of cation exchange capacity (CEC) and the fraction of exchange sites occupied by ‘basic’ cations (predominantly Ca2þ, Mg2þ, Kþ, and Naþ), measurements of decreases in the pH of well-buffered solutions added to the soil sample, and equilibrations of a soil sample with a BaCl2 solution buffered at pH 8.2, followed by titration to pH 5 to determine the quantity of acid that reacted with the buffer. Measures of free CaCO3 content can be used to estimate the total alkalinity of a soil and its buffer capacity against acidification. Diverse chemical reactions occurring on colloid surfaces of clay, oxides, and organic matter are responsible for pH buffering. They involve cation exchange reactions, dissolution and precipitation of sparingly soluble compounds, and surface charge changes in response to pH. In most soils, pH buffer capacity and buffering mechanisms are due to combinations of these reactions, and titrations and quick tests for total acidity do not distinguish between them. None the less, understanding the relative importance of these processes and their chemical basis allows predictions of the buffering behavior of diverse soils, and it allows extrapolation to unstudied soil systems and environments, particularly as related to ecological processes, soil contamination, and soil remediation. Dissolution and Precipitation of CaCO3 and Al(OH)3
Soils containing free CaCO3 have a pH between 7.0 and 9.5 (typically 8.0–8.5) as governed by the
dissolution and precipitation of CaCO3 in water containing partial pressures of CO2 as high as 10 times that of atmospheric levels. Acidity produced by the hydration of dissolved CO2 and by other acid-generating reactions reacts with CaCO3 in accordance with: CaCO3 þ H3 Oþ ¼ Ca2þ þ HCO 3 þ H2 O
½2
In the presence of high levels of soluble or exchangeable Mg2þ or Naþ, soluble MgCO3 or Na2CO3 will form and allow the soil pH to rise to values higher than that controlled by CaCO3–CO2 equilibria (sometimes to pH >10). Addition of base or acid will be buffered by the lime in the soil, thereby maintaining its pH at approximately 8.2. Under strongly acid soil conditions, Al3þ in soil solution and on cation exchange sites hydrolyzes and generates acidity, while the dissolution of Al(OH)3 neutralizes it, leading to large values at pHs <4.5: Al3þ þ 3H2 O ¼ AlðOHÞ3 þ 3Hþ
½3
Soluble and exchangeable Al3þ is released from aluminosilicate clays upon weathering of Si-dominated tetrahedral sheets (releasing soluble monosilicic acid, Si(OH)4), thereby exposing Al-dominated octahedral sheets to Hþ attack. As a result, the buffer capacity of strongly acid soils (pH <5) is dominated by Al chemistry. Cation Exchange Reactions and pH Buffering
In the intermediate pH range between 4.5 and 6.5, cation exchange reactions on permanent aluminosilicate cation exchange sites, and on pH-dependent sites of organic matter and (hydr)oxides govern uptake and release of Hþ and OH as controls on the buffering-capacity size and mechanism for pH control. Permanent, negatively charged cation exchange sites on aluminosilicate clay minerals are due to isomorphous substitution of lower charge cations for Si4þ (e.g., Al3þ) in tetrahedral sheets and for Al3þ (e.g., Mg2þ and Fe2þ) in octahedral positions. Cation exchange sites created in this way during clay mineral formation are dominated by Ca2þ at and above pH 7 and become increasingly satisfied by Hþ and Al3þ at lower pH values. In this way, the so-called base saturation (an intensive variable reflecting the fullness of the acidity samovar) decreases as exchange sites are occupied by Al3þ and Hþ. Organic matter functional groups (principally carboxylic acids, represented by R COOH, and phenw olic acid groups, such as OH on aromatic rings) w are weak acids with association constants (pKas) in the range of 2–7 for carboxylic acids and 7–10
BUFFERING CAPACITY
for phenolic acid groups. These groups deprotonate at pHs >pKa, and protonate at pHs
Oxide and hydroxide coatings of Fe(III), Mn(III,IV), and Al on sand, silt, and clay minerals are dominated by oxygen- and hydroxide-rich planes exposed to soil solution, and changes in pH protonate and
145
deprotonate these functional groups, thereby creating variable or pH-dependent charge colloid surfaces that contribute to buffer capacity: 0 III MeIII OHþ1 2 þ OH ¼ Me wOH þ OH w ¼ MeIII O1 w
½4
where Me(III) represents a trivalent metal structurally associated with the mineral surface, and the charges on the surface become increasingly negative as pH is raised. Highly weathered soils that contain few aluminosilicate clay minerals or organic matter are often dominated by these variable-charge colloids, and they contribute to pH buffering through positive- and negative-charge creation. The relative importance of CaCO3 dissolution and precipitation, Al hydrolysis, cation exchange on permanently charged aluminosilicate minerals, organic-matter functional group protonation and deprotonation, and variation in negative and positive charges on variably charged oxide coatings depend on the relative abundance of these colloids in a given soil. In many soils, more than one of these soil chemical processes are responsible for pH buffering, and titrations of whole-soil materials reflect an integration of the several processes. Table 1 summarizes the relative magnitudes of the buffer capacities associated with these processes.
Table 1 Approximate buffer capacities of soil materials and horizons in the pH range 3.0–10.0
Soil constituent or horizon
Buffering pH range
Buffering index (mmol kg 1 pH unit 1)
Aluminosilicate clays Smectite Vermiculite Illite Kaolinite Soil organic matter
3–10 3–10 3–10 3–10 3–10 5–8
178–333 333–444 44–88 2–11 360–444
Allophane and imogolite
3–10
44–111
Fe(III) and Al(III) (hydr)oxides CaCO3 and MgCO3 Forest floor organic horizons Mineral horizons of agricultural and forest soils
3–10
11–89
>7 3–7
4444 180–360
4–10
53
pH-buffering mechanism
Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Cation exchange; mineral dissolution/precipitation; Al hydrolysis Protonation–deprotonation of weak acid functional groups; conformational changes Protonation–deprotonation of weak acid functional groups; conformational changes Protonation–deprotonation of weak acid functional groups; conformational changes Dissolution/precipitation Metal–ligand complexation/decomplexation; cation exchange; protonation/deprotonation; of weak acid functional groups Cation exchange, mineral dissolution/precipitation, Al hydrolysis, protonation–deprotonation of weak acid functional groups, conformational changes, and metal–ligand complexation/ decomplexation; principally associated with soil organic matter
Source: Bloom PR (2000) Soil pH and pH buffering. In: Sumner M (ed.) Handbook of Soil Science, pp. 333–352. Boca Raton, FL: CRC Press, with permission; McBride MB (1994) Environmental Chemistry of Soils. New York, Oxford University Press, with permission; James BR and Riha SJ (1986) pH buffering in forest soil organic horizons: relevance to acid precipitation. Journal of Environmental Quality 15: 229–234, with permission.
146 BUFFERING CAPACITY
Buffering of Soil Oxidation–Reduction Status The concept of ‘electron activity’ of soils refers to the thermodynamic tendency for electrons to be transferred from reductants (e.g., Fe(II)) to oxidants (e.g., O2), and it is operationally defined by Pt electrode potentials versus a standard reference electrode. When the measured potential is corrected for the reference electrode potential relative to the hydrogen electrode (E0 ¼ 0.0 V), it is designated Eh. This voltage can be converted to pE, a parameter analogous to pH, by dividing Eh in volts by 0.059 (the Nernstian slope factor relating electrode voltage to electron activity). The Eh or pE is the intensity variable that is a measure of electron activity or ‘electron pressure’ in a soil system. Buffering of electron activity is called poise, and is analogous to proton buffering, except that the processes responsible for resistance to change in pE are due to electron transfer reactions, many of which are microbial and enzymatically catalyzed. Heterotrophic microbial respiration in soils uses organic C as the electron source (reductant) and an array of electron acceptors as oxidants. Strict aerobes use dissolved O2 as the required oxidant to derive energy from the oxidation of organic C, the most energy-efficient process for cellular respiration. In flooded soils in which all pores are filled with water, the diffusion rate for dissolved O2 is 10 000 times slower than it is in air. As a result, microbial respiration may deplete available O2 faster than it is replenished, thereby leading to the onset of anaerobic conditions, with a decrease in the measured pE. Under anaerobic conditions, facultative anaerobes and strict anaerobes use alternative electron acceptors, with decreasing metabolic efficiency in the order (depending on pH): NO 3 , Mn(III,IV)(hydr)oxides, Fe(III)(hydr)oxides, SO2 , CO2, and Hþ. 4 The quantities of each of these electron acceptors (the capacity factor for electron activity) in the soil will ‘poise’ the Eh at a given level as governed by the free energy of reaction associated with its reduction, as coupled to the oxidation of organic C. The expected pE and Eh values at which each of these couples would poise the soil are shown at pH 5 and 7 in Table 2. Since the molecular reaction mechanism for many reduction reactions involves the transfer of the electron with a proton to the oxidant (equivalent to the addition of a H atom), the overall reaction raises the soil pH in many cases. As a result, the higher the pH of a soil, the lower the Eh or pE at which a given reduction reaction is expected to occur and at which the soil system will be poised (i.e., at a lower (Hþ), a greater (e), or ‘electron pressure’ is needed to effect the given reduction).
Table 2 Common reduction half-reactions that poise soils at designated pE values at pH 5 and 7; the log K value is the pE for the reaction at pH 0 Poising pE Half-reaction (for 1 electron reduction)
þ
1/4O2 þ e þ H ¼ 1/2H2O þ 1/5NO 3 þ e þ 6/5H ¼ 1/10N2 þ 3/5H2O 1/2Mn3O4 þ e þ 4Hþ ¼ 3/2Mn2þ þ 2H2O Fe(OH)3 þ e þ 3Hþ ¼ Fe2þ þ 3H2O þ 1/8SO2 4 þ e þ 5/4H ¼ 1/8H2S þ 1/2H2O 1/8CO2 þ e þ Hþ ¼ 1/8CH4 þ 1/4H2O e þ Hþ ¼ 1/2H2 1/4CO2 þ e þ Hþ ¼ 1/24C6H12O6 þ 1/4H2O
log K
pH 5
pH 7
20.8 21.1
15.6 14.3
13.6 11.9
30.7
16.7
8.7
15.8 5.2
4.8 1.0
1.2 3.5
2.9 0 0.21
2.1 5 5.9
4.1 7 7.9
Buffering of Ion Activities via Dissolution–Precipitation, Ion Exchange, and Ligand Complexation The soil solution activities of ions other than Hþ and OH are intensity measurements that may be controlled by dissolution–precipitation, ion exchange, and ligand complexation reactions responsible for buffering in soils. Solubility, exchange, and complexation equilibria maintain ion activities in soil solutions as leaching, and plant and microbial uptake; and other biotic and abiotic processes occur and increase or decrease particular ion activities. The disturbance of these equilibria will induce restorative shifts in accordance with the LeChatelier principle, which states that the balance of products and reactants in a dynamic chemical equilibrium will shift in response to a perturbation in order to restore the original balance of reactants and products. For example, if sparingly soluble PbCrO4 is present in a soil, Pb2þ and CrO2 concentrations of approxi4 mately 5.3 107 mol L1 will be maintained in soil solution as controlled by the solubility product, Ksp, of 2.8 1013. Any process that depletes the soluble Pb2þ or CrO2 4 will induce the dissolution of a small amount of solid PbCrO4 to restore the equilibrium activities of the Pb2þ and CrO2 4 . Similarly, additions of these ions to the soil solution will induce precipitation of more PbCrO4. In this way, the ion activities are buffered and maintained in soil solution, as shown in eqn [5]: PbCrO4 ¼ Pb2þ þ CrO2 4
½5
Similar reactions occur via cation and anion exchange on charged colloids to buffer ion activities,
BUFFERING CAPACITY
as discussed above for soil acidity buffering. The suite of exchangeable cations in soils is dominated by Ca2þ, Mg2þ, Kþ, Na2þ, Hþ, and Al3þ, depending on the pH of the soil, the mineralogy of the soil parent material, the aluminosilicate clay mineralogy, plant uptake, microbial processes, and human-induced changes in the soil conditions. Multiple equilibria are established in any soil, but all comprise a capacity factor, represented by the exchangeable cations, and an intensity factor. The capacity factor is sometimes called the ‘quantity’ term, and the quantity-to-intensity ratio (Q/I) is a measure of the nature of the cation or anion exchange buffering reactions responsible for maintaining ion activities in soil solution. Eqn [6] represents a Ca–K cation exchange reaction for such a Q/I relationship: Ex-Ca þ 2Kþ ðaqÞ ¼ Ex-K2 þ Ca2þ ðaqÞ
½6
in which Ex-Ca represents solid-phase, exchangeable Ca, Kþ(aq) is soluble Kþ, Ex-K2 is exchangeable Kþ, and Ca2þ(aq) is soluble Ca2þ. Any environmental conditions that change the activity ratio of Kþ to Ca2þ in soil solution (e.g., changing water content, preferential plant uptake of one cation over the other, precipitation of Ca2þ, or addition of Ca- or Kcontaining materials to the soil) will result in K-forCa or Ca-for-K exchange reactions that restore and buffer the ion activities in soil solution. The balance between the ratios of exchangeable Ca-to-K and soluble Ca-to-K in solution can be described by an equilibrium constant, a measure of the buffer capacity between exchangeable and soluble forms of the cations. The activities of many ions in soil solution are also in equilibrium with complexed or chelated forms of the metals that are also soluble, but in forms in which the positive charge of the cation has been neutralized by negatively charged ligands. Any processes that increase or decrease the activity of the hexaquo form of the ‘free’ form of the cation, e.g., Fe(H2O)3þ 6 , that is in equilibrium with complexed forms, e.g.,
Bulk Density
147
complexed with carboxylic acid groups, hydroxyl ions, or other Lewis bases, will induce a shift in the equilibrium to restore the balance of the complexed and free form of the ion. In this way, the ‘free’ form of the cation (its activity or intensity) is maintained through the equilibrium with the capacity factor, the complexed form. Similar to solubility products and cation exchange equilibria, a stability constant quantifies the relative thermodynamic stability of the complexed and free forms of a metal in a quantity–intensity relationship. See also: Acid Rain and Soil Acidification; Biodiversity; Environmental Monitoring; Eutrophication; Forest Soils; Microbial Processes: Environmental Factors; Nuclear Waste Disposal; Organic Soils; Pollutants: Persistent Organic (POPs); Pollution: Groundwater; Industrial; Remediation of Polluted Soils; Waste Disposal on Land: Liquid; Municipal; Wetlands, Naturally Occurring
Further Reading Bloom PR (2000) Soil pH and pH buffering. In: Sumner M (ed.) Handbook of Soil Science, pp. 333–352. Boca Raton, FL: CRC Press. Brady NC and Weil RR (2002) The Nature and Properties of Soils, 13th edn. Upper Saddle River, NJ: Prentice-Hall. James BR and Bartlett RJ (2002) Redox phenomena. In: Sumner M (ed.) Handbook of Soil Science, pp. 169–194. Boca Raton, FL: CRC Press. James BR and Riha SJ (1986) pH buffering in forest soil organic horizons: relevance to acid precipitation. Journal of Environmental Quality 15: 229–234. Lindsay WL (1979) Chemical Equilibria in Soils. New York: Wiley-Interscience. Marschner H (1995) Mineral Nutrition of Plants, 2nd edn. London, UK: Academic Press. McBride MB (1994) Environmental Chemistry of Soils. New York: Oxford University Press. Reice SR (1994) Nonequilibrium determinants of biological community structure. American Scientist 82: 424–435. Stumm W and Morgan JJ (1996) Aquatic Chemistry, 3rd edn. New York: Wiley-Interscience.
See Porosity and Pore-Size Distribution