Chapter 5. Soil and Soil Solution Chemistry

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  • 5 Soil and Soil Solution Chemistry






    Biogeochemical processes in the terrestrial environment dominate the hydrochem-ical response of small catchments, because streamwater is largely made up ofdrainage water from soils. Biogeochemical processes can be categorized into threemajor groups (Table 5.1.; cf. van Breemen et al., 1983):

    1. Biochemical processes, including interactions between biota and the atmo-sphere (e.g. photosynthesis, respiration, N2fixation), and interactions betweenbiota and soil solution (e.g. assimilation and mineralization).

    2. Geochemical and soil chemical processes, including interactions between solu-tion and the soil solid phase (e.g. cation exchange, adsorption, chemical weath-ering).

    3. Chemical reactions in solution (e.g. hydrolysis, complexation reactions) orbetween solution and atmosphere (e.g. degassing of CO2),

    Processes from all three categories modify the chemical composition of infiltra-tion water. For all major solutes the quantitative importance of individual biogeo-chemical processes has been estimated from input-output budgets and netassimilation rates (e.g. Driscoll and Likens, 1982; van Breemen et aI., 1984;Nilsson, 1985; Binkley and Richter, 1987;Lelong et al., 1988).

    In this chapter we will focus on soil chemical reactions (i.e. categories 2 and 3),and how they may affect concentrations of macro-solutes in streamwater. Besidesa brief presentation of the theory of the dominant soil chemical processes and thespatial and temporal patterns of soil chemical reactions and parameters, someexamples of anthropogenic impacts on soil chemistry and subsequent recovery ofthe soils will be discussed. This chapter will conclude with a method section, deal-ing with sampling and analysis of soils and soil solutions. Aspects of the chemistryof trace metals in soils will be discussed in Chapter 13.

    As indicated in Table 5.1, several biogeochemical processes involve the transferof H+ ions, thus affecting the acid-base chemistry of soils and soil water. Net H+(proton) transfer may be calculated from quantitative estimates of individualchemical processes. By accounting for all proton sources and sinks a proton bud-

    Biogeochemistry of Small Catchments: A Tool for Environmental ResearchEdited by B. Moldan and J.Cerny@ 1994 SCOPE Published by John Wiley & Sons Ltd

    ~ r~')

    ~ l~J

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    Table 5.1 Reaction equations of H+transfer processes and related processes involving biota (after van Breemen et a!., 1983)

    Processes fromleft to right

    Processes fromright to left

    Reaction equation


    H+ -sourceUptake of cationsUptake ofNH/Mineralization + nitrification

    of organic NMineralization +

    oxidation oforganic S

    Mineralization of P

    Dissociation of HzODissociation of COz

    H+- indifferent processesBiota/atmosphereCOz + HzONz + HzO + 2ROHNH3 + ROHHzS + RoOHSOz + RoOH

    H+- transferBiota/solutionM++ROOHNH4 + RoOH

    RNHz + 20z

    RSH + Y,HzO+ Y.OzRHzP04 + HzO

    = CHzO + Oz= 2R 0 NHz + + Y, Oz= RoNHz + HzO= RoSH + HzO= RSH + Y,Oz


    Volatilization of NH3Volatilization of HzS

    H+ - sink=ROOM+W= RNHz + HzO + H+

    Mineralization of M+

    Mineralization of orgo N

    =2o0H + N03- + W Uptake of N03-

    = ROH + sol- + 2H+= ROH + HZP04- + H+

    Uptakeof sol-Uptakeof P

    Solution or solution/atmosphere2HzO = OW + H+COz+ HzO = HC03- + H+

    Protonation of OH-Protonation of HC03-

  • Table 5.1 (continued}

    Processes fromleft to right

    Processesfrom tright to lef

    Reaction equation

    Dissociation oforganic acids

    Complexation of metal ionsL = organic ligand or OH-

    Oxidation of HzSOxidation of SOzNitrification of NH/Nitrification of NOxNitrification ofNz

    Reverse weatheringMn+ IH+ exchangeOxidation of FeZOxidation of FeSDesorption of sol-


    Solution or solution/atmosphere (continued)

    =ROO-+ W

    HL + M+HzS + ZOzSOz + y,Oz + HzO

    NH/ + ZOzNOx + V.(5- Zx)Oz + Y, HzONz + %OZ+ HzO

    Solids/solutionM"+ + n/Z HzOM"+ + nH.exchFez+ + V.Oz+ %HzO

    FeS + Y>i + %HzOexch S04 - + ZHzO

    =ML+H+= sol- + ZH+= sol- + ZH+= N03- + HzO + ZH+=N03-+H+= ZN03- + ZH+

    = nl2MZlnO + nH+= M.exch + nH+= Fe(OHh +ZW= Fe(OHh + SO/- +ZW= exch (OH)z + SO/- +ZH+

    Protonation of organic anions

    Decomplexation of metal ionsSulphate reduction


    WeatheringH+/Mn + exchangeReduction of Fe(OHhReduction of Fe(OHh and SO/-Adsorption of sol-

    Reproduced by permission of Kluwer Academic Publishers.



    getmaybeconstructed.Protonbudgets,whichexpresstherelativeimportanceofall major sources and sinks of acidity, have been used extensively in acidificationresearch (Driscoll and Likens, 1982; van Breemen et aI., 1983). Major sources ofprotons include CO2 dissolution in water, cation assimilation, nitrification andatmospheric acid deposition, while major sinks are cation exchange and chemicalweathering.


    In soil solutionsdissolvedinorganiccarbon(DIC) is abundant,and consists ofH2C03* (C02(aq) + H2C03), HC03- and cOl-. The distribution of DlC speciesin water can be described by equilibrium relationships, where the H2C03* activityis controlled by the partial pressure of CO2in the atmosphere (PC02):

    -Log(H2C03*) = 1.46 - Log(pC02) (5.1)

    where brackets indicate activity (moll-I), and pC02 is in atm.Dissociation of carbonic acid depends on pH and can be described as (Bolt and

    Bruggenwert, 1976):

    -Log(HC03-) = 7.81 - Log (PC02) - pH

    -Log(C032-) = 18.14 - Log(pC02) - 2pH



    For systems open to the atmosphere, pC02 is C. 3xl0-4 atm. However, in soilspC02 (ranging from 10-2 to 10-1 atm; Bolt and Bruggenwert, 1976) is generallyhigher, due to respiration and oxidation of below-ground organic matter.Consequently, DIC concentrations tend to be higher in soil solutions than in sur-face water. Degassing of CO2 is common when soil water emerges (Reuss andJohnson, 1986).

    Carbon dioxide, dissolved in soil water, may react with minerals (includingfeldspars and calcite) according to:

    n12M2/nO+ nC02 + n12H2O+--->Mn++ nHC03-

    generally resulting in soil solution pH values well above 6. Most natural freshwa-ters are in this carbonic acid buffer range (Stumm and Morgan, 1981).

    The presence of dissociated carbonic acid in water gives rise to alkalinity (Alk),where Alk is the equivalent sum of bases that are titratable with strong acid:


    [Alk] = [HC03-] + 2[COl-] + [OH-] - [H+] (5.5)

    with brackets indicating concentrations in moll-I. Alkalinity is also known as theacid neutralizing capacity (ANC). The equivalence point of the acidimetric titra-tion (around pH 4.5) represents an approximate threshold below which most life


    processes in natural waters are seriously impaired. Thus alkalinity is a convenientmeasure for estimating the maximum capacity of a natural water to neutralizeacidity without permitting extreme disturbance of biological activities in the water(Stumm and Morgan, 1981).

    In very dilute natural solutions (e.g. in acidic soils) DIC and therefore [Alk] inEquation (5.5) are low. In such systems additional protolytic systems (of whichhydrolysed Al compounds and natural weak organic acids are the most prominent),may contribute to alkalinity. Assuming that at the equivalence point of the alkalin-ity titration Al is present as AI(OH)2+(Sullivan et at., 1989) the definition of alka-linity becomes:

    [Alk] = [HC03-] + 2[C032-] + [OH-]+ [RCOO-] + 2[AI(OH)z+] + 4[AI(OH)4-] - [H+] (5.6)

    with [RCOO-] representing the concentration of organic anions.Alkalinity is a conservative parameter, i.e. it is pressure and temperature inde-

    pendent. For example, degassing of CO2 results in the removal of equivalentamounts of H+ and HC03- from solution, thus causing no change in [Alk] inEquations (5.5) and (5.6). By contrast, degassing of CO2 may result in a signifi-cant increase in solution pH, particularly in solutions with positive alkalinity(Reuss and Johnson, 1986; Suarez, 1987). Degassing in solutions with a negativealkalinity (i.e. solutions where strong mineral acids dominate) will have little or noeffect on pH, however.


    Soil organic matter (SOM) can be subdivided into non-humified and humifiedmaterial. Non-humified substances are not or are only slightly altered after decayof tissue from living organisms and include, e.g. carbohydrates, amino acids, pro-tein, lignin, hormones and low molecular weight organic acids (Tan, 1986).Humified substances are decomposition products of non-humified constituents andinclude complex compounds such as humin, fulvic acid (FA), hymatomelanic acid,humic acid (HA) and their hydroxybenzoic acid derivatives (Tan, 1986).

    The concentrations of non-humified organic acids are generally low and manyof these acids can only be detected by thin layer or gas chromatography.Nevertheless, with their rapid turnover, low molecular weight organic acids mayplaya significant role in mineral weathering. In most soils the contents of HA andFA are considerably higher than those of the non-humified organic acids. Themajor reason for the importance of FA and HA in soil chemistry is the presenceand position of functional groups (particularly carboxyl and phenolic hydroxylgroups), which make FA and HA effective in cation exchange and complexationreactions (Tan, 1986). Charge characteristics of humic substances depend upon theextent of dissociation of the functional groups. At pH < 3 HA and FA behave as


    uncharged polymers, whereas at pH > 3 they behave as negatively charged poly-electrolytes, due to dissociation of carboxyl groups (3 < pH < 9) and phenolichydroxyl groups (pH> 9) (Tate and Theng, 1980).

    Fulvic and humic acids may form metal complexes of high stability throughchelation, where metal ions are bound as bidentate e.g. by carboxyl and phenolichydroxyl groups. Chelation promotes the dissolution of metals from soil minerals.For example, in podzols chelation of Al and Fe in the surface mineral soil and sub-sequent leaching eventually result in Al and Fe depleted eluvial (E) horizons. Thesolubility of chelates decreases with increasing sesquioxide content, however, thusresulting in precipitation in the illuvial (podzol B) horizon.

    In soils, DOC concentrations decrease sharply from several tens of mg r1 in the0 and E horizons to only a few mg r1 in the mineral soil (e.g. McDowell andWood, 1984; Cronan and Aiken, 1985). Besides physicochemical sorption and pre-cipitation, this decrease in DOC is likely due to mineralization.

    As indicated in Section 5.1.2, DOC may contribute to the alkalinity of soil solu-tions and surface waters. On the average, organic acids in natural water contribute5 to 10 J.1eqof anionic charge per mg DOC. To estimate the anionic charge associ-ated with DOC in individual samples the empirical model by Oliver et ai. (1983)has frequently been used (e.g. Driscoll et ai., 1989a). Alternatively, Cronan andAiken (1985) and Schecher and Driscoll (1987) used analogues of monoprotic,diprotic and triprotic acids with optimized dissociation constants. Recently,Tipping et al. (1988) proposed a deterministic model which treats organic acids asmacroions, where interactions are modelled in terms of discrete binding sites. Thismodel accounts for electrostatic (macroionic) effects which depend upon molecu-lar charge (determined by the bound cations) and ionic strength.

    Because soil organic matter is a major source of negative charge in many soils itacts as an important cation exchanger, which may represent a significant chemicalbuffer (e.g. the forest floor; James and Riha, 1986). In addition, soil organic mattermay complex trace metals (e.g. AI), thus reducing phytotoxic effects (e.g. Bloomet ai., 1979). Soil organic matter is also important as a store for Nand S, whichmay be liberated slowly upon decay.


    In many soils the solid phase carries a net negative surface charge. For clay miner-als this charge is a result of isomorphous substitution, where structural cations ofhigher valence are replaced by cations of lower valence (e.g. Si4+is replaced byAI3+).As a result, clay minerals have a permanent negative surface charge.

    Oxides and SOM also have charged surfaces: With increasing pH (i.e. increas-ing activity of OH- ions) H+ is dissociated from oxide surfaces or from organicfunctional groups, thus resulting in a negative charge. Because dissociationincreases with pH and ionic strength, such charge is termed variable charge.

    At low pH values oxides may bind H+, which results in a positive surfacecharge. The pH values at which positively charged groups quantitatively equal


    negatively charged groups (i.e. the net surface charge is zero) are called zero pointof charge (ZPC).

    In soils the overall electroneutrality is maintained by an excess of electrostati-cally attracted counterions in proximity to the charged surface. In the case of nega-tively charged surfaces a diffuse double layer will result where cations are inexcess of anions. The excess ions, termed exchangeable cations (anions) may, bydefinition, be exchanged with neutral salts. The quantity of exchangeable cations(in meq kg-1 dry soil) is defined as the cation exchange capacity (CEC). Similarlythe total amount of exchangeable anions is defined as the anion exchange capacity(ABC). In most soils the AEC is quantitatively far less important than the CEC.

    Exchangeable cations are available to plants, for example through exchangewith H+ liberated by the roots. Exchange reactions are also responsible for theretention of freshly introduced cations into the soil solution. In this way the CECgives the soil a buffering capacity, which may slow down the leaching of nutrientcations and positively charged pollutants.

    Since cation exchange is a relatively fast process it modifies the chemical com-position of infiltration water. The exchange reaction of cations M and N, with acharge of m+and n+,respectively, may be represented as:

    Ads-Nm + nMm+ ~ Ads-Mn + mNn+ (5.7)

    Mathematically this reaction has been described by the Gapon equation:

    {Ads(M)/Ads(N)} = KNMgx {(Mm+)1/m/(Nn+)I/n} (5.8)

    with the left-hand side being the ratio of adsorbed Mover N (both in meq per massunit exchanger). The right-hand side contains the reduced activity ratio, where thecation activities are raised to a power equal to the reciprocal of their valence.KNMgin this equation represents the (Gapon) selectivity constant, which shouldbe constant over a wide range of conditions. The selectivity constant differs fromone, because small size (hydrated) ions are generally preferred over large ones(high KNM),due to their smaller distance of approach of the charged surface. Alsothe surface structure of the exchanger may affect the selectivity constant (e.g. incase of a porous exchanger, where the binding may be determined by the "naked"cation; Bolt and Bruggenwert, 1976).

    Alternative mathematical descripti...


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