Environmental Soil Chemistry - Chapter 8 doc

of Soils Oxidation–Reduction Reactions and Potentials electron transfer.. Oxidation occurs if there is a loss of electrons in thetransfer process while reduction occurs if there is a gai

of Soils

Oxidation–Reduction Reactions and Potentials

electron transfer Oxidation occurs if there is a loss of electrons in thetransfer process while reduction occurs if there is a gain of electrons.The oxidized component or oxidant is the electron acceptor and the reducedcomponent or reductant is the electron donor Table 8.1 lists oxidants and reductants found in natural environments The electrons are not free inthe soil solution; thus the oxidant must be in close contact with thereductant Both oxidation and reduction must be considered to completelydescribe oxidation–reduction (redox) reactions (Bartlett and James, 1993;

Patrick et al., 1996).

To determine if a particular reaction will occur (i.e., the Gibbs freeenergy for the reaction, ΔGr<0), one can write reduction and oxidation half-reactions (a half-reaction or half-cell reaction can be referred to as a redoxcouple) and calculate equilibrium constants for the half-reactions Redoxreactions of soil oxidants can be defined conventionally by the general half-

reduction reaction (Patrick et al., 1996)

Ox + mH++ ne–→ Red, (8.1)where Ox is the oxidized component or the electron acceptor, Red is the

reduced component or electron donor, m is the number of hydrogen ions participating in the reaction, and n is the number of electrons involved in the

reaction The electrons in Eq (8.1) must be supplied by an accompanyingoxidation half-reaction For example, in soils, soil organic matter is theprimary source of electrons Thus, to completely describe a redox reaction,

an oxidation reaction must balance the reduction reaction

Let us illustrate these concepts for the redox reaction of Fe(OH)3

reduction (Patrick et al., 1996):

4Fe(OH)3+ 12H++ 4e–→ 4 Fe2++ 12H2O (reduction) (8.2)

4 Fe(OH)3+ CH2O + 8H+→ 4 Fe2++ CO2+ 11H2O (net reaction), (8.4)where CH2O is soil organic matter Equation (8.2) represents the reductionhalf-reaction and Eq (8.3) represents the oxidation half-reaction The reduction (Eq (8.2)) reaction can also be described by calculating ΔGr,the Gibbs free energy for the reaction,

where ΔGrois the standard free energy change for the reaction The Nernstequation can be employed to express the reduction reaction in terms ofelectrochemical energy (millivolts) using the expression ΔGr = –nFE such that (Patrick et al., 1996)

where Eh is the electrode potential, or in the case of the reduction

half-reaction in Eq (8.2), a reduction potential, E° is the standard half-half-reaction

reduction potential (with each half-reaction, for example, Eqs (8.2) and(8.3), there is a standard potential; the standard potential means the activities

of all reactants and products are unity), F is the Faraday constant, n is the number of electrons exchanged in the half-cell reaction, m is the number of

protons exchanged, and the activities of the oxidized and reduced species are

in parentheses Determination of Eh will provide quantitative information

on electron availability and can be either an oxidation or reduction potentialdepending on how the reaction is written (see Eqs (8.2)–(8.3)) Oxidationpotentials are more often used in chemistry, while in soil chemistry reductionpotentials are more frequently used to describe soil and other natural systems

(Patrick et al., 1996) It should also be pointed out that the Nernst equation

is valid for predicting the activity of oxidized and reduced species only if thesystem is at equilibrium, which is seldom the case for soils and sediments

As noted in Chapter 7, the heterogeneity of soils which promotes transportprocesses causes many soil chemical reactions to be very slow Thus, it isdifficult to use Eh values to quantitatively measure the activities of oxidizedand reduced species for such heterogeneous systems (Bohn, 1968)

Oxidation–Reduction Reactions and Potentials 247

TABLE 8.1. Selected Reduction Half-Reactions Pertinent to Soil, Natural Water, Plant, and Microbial Systems a

TABLE 8.1. Selected Reduction Half-Reactions Pertinent to Soil, Natural Water, Plant, and Microbial Systems a (contd)

aFrom Bartlett and James (1993), with permission.

bCalculated for reaction as written according to Eq (8.14) Free energy of formation data were taken from Lindsay (1979) as a primary source, and when not available from that source, from Garrels and Christ (1965) and Loach (1976).

c Calculated using tabulated log K°values, reductant and oxidant = 10 –4 M soluble ions and molecules, and activities of solid phases = 1; partial pressures for gases that are pertinent to soils: 1.01 × 10 –4 MPa for trace gases, 2.12 × 10 –2 MPa for O2, 7.78 × 10 –2 MPa for N2, and 3.23 × 10 –5 MPa for CO2.

dValues not listed by Loach (1976).

Using the values of 8.31 J K–1mol–1for R, 9.65 × 104C mol–1for F, and

298 K for T and the relationship ln(x) = 2.303 log(x), Eq (8.6) becomes

It is obvious from Eqs (8.6)–(8.7) that Eh increases as the activity of theoxidized species increases, decreases with increases in the activity of thereduced species, and increases as H+activity increases or pH decreases If the

ratio of protons to electrons is 1 (i.e., m/n = 1), one would predict that Eh

would change by 59 mV for every unit change in pH Thus, one couldpredict the Eh at various pH values by using the 59-mV factor However, this relationship assumes that redox controls the pH of the system Thisassumption is valid for solutions, but in soils pH buffering is affected by soilcomponents such as silicates, carbonates, and oxides, which are not involved

in redox reactions Thus, it may be inappropriate to apply the 59-mV factor

(Patrick et al., 1996).

Eh is positive and high in strongly oxidizing systems while it is negativeand low in strongly reducing systems (There is not a neutral point, as oneobserves with pH.) Eh, like pH, is an intensity factor The oxygen–nitrogenrange has been defined by Eh values of +250 to +100 mV, the iron range as+100 to 0.0 mV, the sulfate range as 0.0 to –200 mV, and the methane–hydrogen range as <200 mV (Liu and Narasimhan, 1989)

Eh vs pH and pe vs pH Diagrams

Diagrams of the activities of Eh vs pH can be very useful in delineating theredox status of a system Figure 8.1 shows such a diagram for soils The pHrange was narrower in reduced soils (negative Eh) than in oxidized soils

(positive Eh) Based on these results, Baas Becking et al (1960) divided the

soils into three categories: normal (oxidized), wet (seasonally saturated), and waterlogged (semipermanently saturated) (Fig 8.1)

The reduction half-reaction given in Eq (8.1) can also be expressed in

terms of an equilibrium constant K ° (Patrick et al., 1996).

Expressed in log form Eq (8.8) becomes

log K ° = log(Red) – log(Ox) – nlog (e–) – mlog (H+) (8.9)

The –log (e–) term in Eq (8.9) is defined as pe in a similar way as pH is

expressed as –log (H+) The pe is an intensity factor as it is an index of the

electron free energy level per mole of electrons (Ponnamperuma, 1972)

Thus, pe and pH are master variables of a soil and must be known to

completely understand the equilibrium state of a soil Moreover, to fully

determine the redox status of a soil, pe and pH cannot be separated (Bartlett and James, 1993) In strongly oxidizing systems the e–activity is low and pe

is large and positive In reducing systems pe is small and negative Sposito (1989) proposed “oxic” (oxidized) soils as those with pe >7, “suboxic” soils in

the pe range between +2 and +7, and “anoxic” (reduced) soils with pe <+2,

all at pH 7 These ranges are consistent with redox control by oxygen–nitrogen, manganese–iron, and sulfur couples (James and Bartlett, 2000)

The pe range of most soils is –6 to +12 (Lindsay, 1979) Rearranging

Eq (8.9) one arrives at an expression that relates pe to pH,

pe = [(log K ° – log (Red) + log (Ox))/n] – m/n pH, (8.10)

which represents a straight line with a slope of m/n and an intercept given in brackets The intercept is a function of log K° for the half-reaction and theactivities of the oxidized and reduced species When there is a one-electron

transfer (i.e., n = 1) and consumption of one proton (i.e., m = 1), and when

(Red) = (Ox), Eq (8.10) is simplified to

where 5.71 is derived from the product of (RT)(2.303), R is (0.008314 kJ mol–1

K–1), and T = 298.15 K Therefore, log K° could be estimated by knowingthe free energies of formation (ΔGfo) of H2O and the Red and Ox speciessince those for H+and e–are zero by convention (Bartlett and James, 1993)

Information in Box 8.1 shows how one would calculate log K ° and pe

for a reduction half-reaction at pH 5 and 7 using Eqs (8.11)–(8.14)

The values of log K° can be used to predict whether a reduction and

oxidation reaction will combine to effect the transfer of electrons fromreductant to oxidant Table 8.1 lists a number of reduction half-reactions

important in natural systems The log K° values are given in descending order

and are pe values at pH 0, when the activities of oxidant and reductant are 1, and are standard reference pe values for the reactions The larger the values

FIGURE 8.1. Eh–pH characteristics of soils.

From Baas Becking et al (1960), with permission from the University of Chicago Press.

Oxidation–Reduction Reactions and Potentials 251

of log K ° or pe, the greater the tendency for an oxidant (left side of the

reaction equation) to be reduced (converted to the right side of the reaction equation) Therefore, an oxidant in a given reduction half-reaction

half-can oxidize the reductant in another half-reaction with a lower pe, at a

particular pH As an example, Mn(III,IV) oxides could oxidize Cr(III) to

Cr(VI) at pH 5 because the range of pe values for reduction of Mn

(12.8–16.7) is higher than that for Cr(VI) reduction (10.9) (Bartlett andJames, 1993) In field moist soils over a pH range of 4–7 it has indeed beenobserved that Mn(III,IV) oxides can oxidize Cr(III) to Cr(VI) (Bartlett andJames, 1979; James and Bartlett, 1983)

The pe–pH relationship expressed in Eq (8.10) can be used to

determine whether an oxidation–reduction reaction can occur spontaneously,i.e.,ΔGr< 0 Figure 8.2 shows pe vs pH stability lines between oxidized and

reduced species for several redox couples If thermodynamic equilibrium is

present, the oxidized form of the couple would be preferred if the pe and pH

region was above a given line and the reduced form would be favored below

a given line (Bartlett, 1986) The line for Fe is often considered the dividingpoint between an aerobic (oxidized) and an anaerobic (reduced) soil

In aerobic soils oxidized species stay oxidized even though the

thermodynamic tendency is toward reduction, as indicated by the high pe.

Below the iron line, reduced species are prevalent, even though thethermodynamic tendency is toward oxidation Sulfide is easily oxidized andnitrite is easily reduced (Bartlett and James, 1993)

BOX 8.1 Calculation of log K° and pe

The reduction half-reaction below (see Table 8.1) shows the reduction of

Fe3+to Fe2+,

In this reaction there is one electron transfer, i.e., n in Eq (8.8) is 1, there is consumption of one proton, i.e., m = 1, and (Fe3+) = (Fe2+) is animposed condition Thus Eq (8.10) reduces to Eq (8.11) and at pH 0,

Eq (8.12) results Relating ln K ° to ΔGro, one can employ Eq (8.13),

CO2(2.0x10-3

MPa) glucose (.01M

)

.

2-

To calculate pe at pH 5 and pH 7, one would use Eq (8.11) For

pH 5 and substituting in the value of 15.2 for log K°,

Figure 8.3 illustrates a pe–pH diagram for several Mn species One sees

that Mn oxides can oxidize Pu(III) to Pu(IV), V(III) to V(V), As(III) to

As(V), Se(IV) to Se(VI), and Cr(III) to Cr(VI), because the pe for each of these couples is below the pe for Mn oxides It has been shown that Mn

oxides in soils can indeed effect oxidation of Pu(III), As(III), Se(IV), and, asnoted earlier, Cr(III) (Bartlett and James, 1979; Bartlett, 1981; Amacher and

Baker, 1982; Moore et al., 1990) The environmental aspects of some of these

oxidation processes are discussed later in this chapter

Another term often used in studying redox chemistry of soils is poise.The poise of a redox system is the resistance to change in redox potential withthe addition of small amounts of oxidant or reductant Poise increases withthe total concentration of oxidant plus reductant, and for a fixed totalconcentration it reaches a maximum when the ratio of oxidant to reductant

is 1 (Ponnamperuma, 1955)

Measurement and Use of Redox Potentials

Measurement of redox potentials in soils is usually done with a platinumelectrode This electrode will transfer electrons to or from the medium, but

it should not react with the medium Once the platinum electrode iscombined with a half-cell of known potential, reducing systems will transferelectrons to the electrode while oxidizing systems will remove electrons fromthe electrode When experimental redox potential measurements are done,there is no electron flow and the potential between the half-cell composed ofthe platinum in contact with the medium and the known potential of thereference electrode half-cell is determined with a meter that reacts to the

electromotive force or potential (Patrick et al., 1996) A number of investigators

have noted that measurement of redox potentials in aerated soils is questionabledue to the lack of poising of reduction–oxidation systems that are well aeratedand have plentiful quantities of oxygen (Ponnamperuma, 1955) Soil atmosphericoxygen measurements are preferred to characterize well-aerated soils Thus,redox measurements are most reliable for flooded soils and sediments

35 25 15 5 -5 -15

FIGURE 8.3. A pe–pH diagram for Mn 3+ , MnO 2 , and Mn 3 O 4 ;

as compared with reduction between pH 5 and 7 for Co, Cr, Se, As,

V, and Pu Activity for ionic species is 10 –4 M From Bartlett and

James (1993), with permission.

Redox potentials can be very useful in characterizing the oxidation–reduction status of a soil Oxidized soils have redox potentials of +400 to+700 mV Seasonally saturated soils have redox potentials of +400 to +700

mV (oxidized) to highly reduced (–250 to –300 mV) (Patrick et al., 1996).

Redox potentials can help one predict when reducing conditions will begindue to depletion of oxidants such as oxygen and nitrate, and the initiation ofoxidizing conditions when oxygen is reintroduced in the soil Redoxpotentials can also provide information on conditions that are favorable for

increased bioavailability of heavy metals (Gambrell et al., 1977; Reddy and Patrick, 1977), changes in plant metabolism (Mendelssohn et al., 1981), distribution of plant species (Josselyn et al., 1990), and location of wetlands (Faulkner et al., 1989).

If redox potential data are combined with other information such asdepth to the water table and oxygen content of the soil, even more accurateinformation can be gleaned about the wetness of an environment In nonwetlandenvironments the Eh and oxygen content do not change much during theyear Transitional areas may be either oxidized or reduced as the water tablerises and falls The redox potentials are low until after the water is drainedand oxygen moves through the soil Wetland sites that have low redox have

had long periods of flooding and soil saturation (Patrick et al., 1996).

Redox data are also useful in understanding the morphology and genesis

of the soil The color of a soil and the degree of mottling (spots or blotches

of different colors or shades of color interspersed with the dominant color;Glossary of Soil Science Terms, 1997) can reveal much about the soil’smoisture status Both color and mottling depend on the redox chemistry of

Fe in the soil When the soil is saturated for long times Fe oxides are reducedunder low redox potentials, and the soil will exhibit a gray color Soils thatundergo alternate oxidation and reduction cycles are usually mottled (Patrick

acceptors, temperature, and period of submergence Native or added SOMenhances the first Eh minimum, while nitrate causes the minimum todisappear Temperatures above and below 298 K slow the Eh decrease but theretardation varies with the soil It is greatest in acid soils and not observable

in neutral soils with high SOM

Redox Reactions Involving Inorganic and Organic Pollutants 255

Ponnamperuma and Castro (1964) and Ponnamperuma (1965) havelisted the effects of soil properties on Eh changes in submerged soils as follows:soils high in nitrate, with more than 275 mg kg–1 NO3–, exhibit positivepotentials for several weeks after submergence; soils low in SOM (<1.5%) orhigh in Mn (>0.2%) have positive potentials even 6 months after submergence;soils low in active Mn and Fe (sandy soils) with more than 3% SOM reach

Eh values of –0.2 to –0.3 V within 2 weeks of submergence; and stablepotentials of 0.2 to –0.3 V are attained after several weeks of submergence

Redox Reactions Involving Inorganic and Organic Pollutants

As noted in Chapter 2, metal oxides/hydroxides, e.g., Fe(III) andMn(III/IV), are quite common in soils and sediments as suspended particlesand as coatings on clay mineral surfaces Manganese(III/IV), Fe(III), Co(III),and Pb(IV) oxides/hydroxides are thermodynamically stable in oxygenatedsystems at neutral pH However, under anoxic conditions, reductivedissolution of oxides/hydroxides by reducing agents occurs as shown below

Mn(III)OOH(s) + 3H++ e–= Mn2++ 2H2O E° = + 1.50 V (8.15)Mn(IV)O2(s) + 4H++ 2e–= Mn2++ 2H2O E° = + 1.23 V (8.16)Changes in the oxidation state of the metals associated with the oxides abovecan greatly affect their solubility and mobility in soil and aqueousenvironments The reductants can be either inorganic or organic

There are a number of natural and xenobiotic organic functional groupsthat are good reducers of oxides and hydroxides These include carboxyl,carbonyl, phenolic, and alcoholic functional groups of SOM Microorganisms

in soils and sediments are also examples of organic reductants Stone (1987a)showed that oxalate and pyruvate, two microbial metabolites, could reduceand dissolve Mn(III/IV) oxide particles Inorganic reductants include As(III),Cr(III), and Pu(III)

Table 8.2 gives standard reduction potentials (E°) for oxide/hydroxide

minerals at a metal concentration of 1.0 M and potentials (E°′) determinedunder more normal environmental conditions of pH 7 and a metalconcentration of 1 × 10–6 M Oxidant strength decreases in the orderMn(III,IV) oxides > Co(III) oxides > Fe(III) oxides The Fe oxides are moredifficult to reduce than Mn(III,IV) oxides Thus Fe(II) is easier to oxidizethan Mn(II) Reduction potentials for oxidation of some important organichalf-reduction reactions are provided in Table 8.3

If the potential of the oxidant half-reaction is higher than the reductanthalf-reaction potential, then the overall reaction is thermodynamically

favored Thus, comparing E°′values in Tables 8.2 and 8.3, one observes that