Tải bản đầy đủ - 0 (trang)
VIII. Free Radicals in Redox Processes

VIII. Free Radicals in Redox Processes

Tải bản đầy đủ - 0trang



+ 3e- + 3H+ H,O + OH (hydroxyl free radical)

0, + 4e- + 4H+

2H,O (water)

HO; + H,O,

H,O + 0, + OH.




OH. + H 2 0 2 - H 2 0 + H Q






2H202 a 2 H 2 0 + 0,






2HO; x H Z O Z 0,




+ 3H+


+ Mn3? + 2H,O





With one electron, a hydrogen atom is a free radical, a single protonated

electron in frantic search for an electron mate. It is the simplest and most

reactive (least stable) free radical. As H,, however, it is quite stable, because the electrons are paired.

Catalases or fresh recently oxidized Mn oxides at pH > 6 (Bartlett,

I98 la) will catalyze the oxidation of one of the oxygens in H,Oz by the

other and thereby will destroy the peroxide by dismutation [Eq. (34)]. In

acid media, oxidized Mn will oxidize H,O,. The dismutation of HO; to

H,O, and 0, by superoxide dismutase (SOD) enzymes, Eq. (33), followed

by H,O, dismutation to 0, and H,O or oxidation of H,O, to O,, make

aerobic life possible (Fridovich, 1975; Halliwell, 1974) by preventing the

formation of the biodestructive hydroxyl free radical by Eq. (3 1). Reforming HO; by Eq. (32) is prevented.

Oxygen free radicals are much more reactive than 0,. Probably free

radical mechanisms explain why kinetically very slow and seemingly unlikely redox transformations sometimes occur readily. The hydroxyl free

radical, (OH the superoxide free radical, -OF, and the supermanganese

free radical ( Mn33are close to being the most powerful oxidizing agents in

soil systems, and superoxide and supermanganese also are the most powerful reducing agents (Table I, Fig. 3, and see Sections VI,A and VI,C)

(Bartlett, 198 I a).

Oxygen free radicals are among the few species having the thermodynamic capability for oxidizing Mn(I1). This means that Mn is one of the

few elements that is capable of scavenging these radicals and protecting life




forms from their biotoxicity. By scavenging free radicals, Mn disrupts the

tendency toward thermodynamic equilibrium between 0, and soil organic

matter (or living roots), allowing the persistence of metastable humus and

roots in an oxidative environment.

Equation (35) shows the MnO, oxidation of hydroquinone by a single

electron step to form two free radicals, the semiquinone free radical and

the supermanganese free radical, Mn? It seems that free radicals could be

the latchkeys to the linking together of reducing phenolic compounds into

organic polymers that are stable in the presence of 0,. For example, a

single OH free radical initiates the linking process in the formation of a

polyethylene chain. The free radical attacks and breaks the 71 bond of an

ethylene molecule, forming a new free radical, which then attacks another

ethylene molecule and forms another new free radical, and so on, until

thousands of molecules are joined (Zumdahl, 1986).





The redox elements, oxygen and carbon, are basic to the building and

functioning of soils and all living systems, and basic to the functioning of

oxygen and carbon in life’s overall redox scheme are Mn and Fe (Figs. 7

and 8). As the key that unlocks oxygen from water in the process of

photosynthesis, Mn is responsible for the presence of the oxygen in the

atmosphere of the planet Earth. Soil Mn also is a protector of life, the

scavenger of death-dealing oxygen free radicals.

Manganese and iron together provide the key to the establishment of the

organic mantle, the humified soil top layer that covers the surface of the

earth and serves as the nurturing home for the roots of all plants and

carbon-recycling microorganisms. It is the sole provider of food for a

variety of Earth’s creatures, including people.

Iron is vital in the life systems of all plants, animals, microbes, and soils.

In biological systems, Fe appears to play many roles similar to those of Mn.

In animals, on the other hand, Fe is supreme, and Mn, although essential,

displays some toxicity tendencies that may prevent its predomination of

Fe. Iron is the chief camer of oxygen in blood, and it is a camer of

electrons to oxygen in both plants and animals. Except in the soil redox

scheme, Fe appears to have a bigger role than Mn in many living redox

systems. In redox of soils, Fe plays second fiddle to Mn, but the melody of

soil processes requires both metals.



H20 /02

0 f r e e radical





everse dismutatio





0 free radicals

oxidative polymerization




Figure 7. The manganese redox system: rather than a simple redox cycle or even a series

of cycles, the Mn system appears to be more a complex web of interacting redox transformations.

Fe(I1) in Primary Minerals



[Low Solubility Fe(lI1) Minerals]


chelat ion

acid dissolution

02-Carrying Proteins


0 free radicals

Organic f r e e radicals

Hemoglobin- Fe ( II)

Met hemoglo bin- Fe( III)

Reduced N

Cr(II1. VI)

Cytochromes-Fe( 11,111)



Increases in:

pe, pH, darkness, H




Catalases-Fe( 111)

Per o x i d a s e s - Fe( III)

Figure 8. Iron redox cycling: though not depicted here as a circular scheme, the reversibility of the various Fe reactions shown here indicates that Fe is involved in cycling, once it

moves out of its primary mineral origins.







Direct oxidation of soil organic compounds by atmospheric 0, is a rare

occurrence. The majority of apparently spontaneous oxidation reactions

are catalyzed by microbial enzymes, metal oxides, peroxides, or free radicals. Such species serve as electron camers, that is, substances that will

oxidize reduced substances and will in turn be oxidized by more highly

oxidized species. Proteins containing metals-usually Fe, Cu, or Mnserve as redox enzymes in microbial systems carrying electrons from reduced carbon to 0,. These enzymes are essential links in the respiration of

all aerobic organisms. Flavoproteins without metals can transfer electrons

to O,, but the rate is too slow to account for aerobic respiration.

The Fe in hemoglobin must be in the ferrous form in order for it to

combine reversibly with O2and act as the all important carrier of 0, in the

blood of animals (Fruton and Simmonds, 1961). If the Fe is oxidized to

Fe( III), methemoglobin is formed, which does not combine with 0,. The

danger of nitrate in drinking water, spinach, or forage plants is that nitrite,

formed from nitrate by reduction in an infant’s intestine, or in the rumen

of a cow, can oxidize hemoglobin to methemoglobin, causing methemoglobinemia, the inability of the blood to carry oxygen. In the heme structure

of cytochrome c, the Fe is reversibly oxidized and reduced, but in catalases

and peroxidases, also heme compounds, the Fe remains trivalent.

In the enzyme cytochrome oxidase, Fe(II1) is reduced by an organic

compound that is oxidized in the process. The Fe(I1) formed then

“cames” an electron to atmospheric 0,, which oxidizes it spontaneously

back to Fe(II1). Names of redox enzymes are sometimes confusing. An

enzyme carrying an electron from an electron donor to O2 is called an

oxidase because the donor is oxidized. But an enzyme that oxidizes a

carbon compound by transfemng an electron to an oxidized substance

other than O2 or peroxide usually will be referred to as a reductase. An

oxidase that removes both a proton and an electron from an electron-donating substance is called a dehydrogenase, because a proton and an electron

comprise a hydrogen atom. Peroxidase is a peroxide reductase (pe = -4.6

at pH 7; Table I ) usually refined from horseradish. A molecule (about

40,000 g mol-’) contains one atom of Fe(I1) (Fruton and Simmonds,

196 I). To pass along an electron to H,O,, the Fe(II1) borrows an electron

from a carbon atom to form, for an instant, an atom of Fe(I1). The Fe( 111)

is restored during that same instant as the H 2 0 2is reduced.

Because it is so readily oxidized by O,, Fe( 11) can catalyze the oxidation

of a phenolic compound that complexes Fe( 111). The phenolic is oxidized

to a quinone when it reduces the Fe( 111) to Fe( II), which is reoxidized by

0, (Stumm and Morgan, 1981). The Fe(II1) formed is complexed and



reduced by the excess phenolic compound, and so on. Trace levels of Fe or

other metals can catalyze the oxidation and spoilage of foods by similar

mechanisms. Citric acid is added to prepared food to tie up the Fe( 111) and

prevent its complexation and reduction by the easily oxidized electrondonating food. This reaction occurs commonly in soils in the absence of

light. More difficult oxidations that apparently do not occur in darkness

frequently do take place when the Fe and reduced carbon are exposed to

energy from sunlight (see Section XI).

Because its reoxidation is so much more difficult than that of Fe, Mn is

less effective than Fe in catalyzing complete oxidation of organics. Manganese is more likely to oxidize organic residues partially (making free radicals), setting them up for further microbial breakdown. It is interesting to

note that of the three trivalent redox cations, Cr(II1) can hydrolyze or it

can oxidize, Mn( 111) can only oxidize, and Fe( 111) can only hydrolyze.





1 . Ultimate Electron Acceptor

The outstanding contribution of manganese to life on Earth appears as a

simple entry in the upper right-hand comer of Fig. 7. Both Mn(II1) and

Mn(IV) are powerful oxidants in the soil redox system, especially the free

radical supermanganese Mnw ion, which has the thermodynamic capability of oxidizing 202- to 0, gas. It does this in the leaves of green plants

exposed to radiation from the sun in the process of photosynthesis and

thereby is responsible for creating the oxygen in the atmosphere. However,

the mechanism for this profound redox transformation is only partially

understood (Brudvig and Crabtree, 1989; Thorp and Brudvig, 1991).

It is possible that a singlet chlorine free radical is the direct electron

acceptor, and that the Mn role is one of accepting an electron from C1- to

form (CIS). Marschner (1986) discusses the evidence that chloride acts as a

cofactor in the Mn-containing 0,-evolving system.

2. Proportionation and Disproportionation

Mn2+ is a soluble or exchangeable cation. It forms by reduction of

Mn(II1) or Mn(1V) by a multiplicity of easily oxidizable, reduced organics,

often microbial by-products. The reductive dissolution of Mn oxides by

microbial metabolites has been studied extensively by Stone and Morgan

( 1984). Simultaneous formation of Mn( 11) and Mn( IV) can take place by

the thermodynamically spontaneous disproportionation, or dismutation,



of two Mn( 111) ions. One Mn( 111) loses an electron to the other to become

Mn(IV), while the electron-accepting Mn( 111) forms Mn( 11). In a reverse of

this dismutation reaction, two molecules of free radical Mn( 111) are constituted when a Mn( 11) gives up an electron to Mn( IV) as follows:


Mn2+ MnO,

+ 3(COOH),



+ 2H20 + 2H+


This reverse dismutation would not be thermodynamically spontaneous

if it were not for the energy of formation of the organic acid-Mn(II1)

complex. An organic acid that readily couples with Mn(II1) and also is

easily oxidized by it can drive the reverse dismutation equation to the right.

With low-molecular-weight organic acids, the Mn( 111) complexes are soluble and range in color from yellow to yellowish brown to red. To be readily

oxidized by Mn( III), the acid must have an oxygen on a carbon adjacent to

a carboxyl group. Oxalic, citric, and tartaric acids are examples, but succinic acid, which chelates Fe( 111) and Al( III), does not complex Mn( 111)

and will not drive the redox.

Pyrophosphate is another ligand that drives this redox reaction by binding strongly to Mn(II1) to form a violet-pink color (Dion and Mann,

1946). Loss of complex color accompanying the oxidation of dissolved

organic carbon by the Mn(II1) serves as a simple colorimetric method for

measuring dissolved organic carbon (Bartlett and Ross, 1988). Pyrophosphate-bound Mn(II1) is not as powerful an oxidizing agent as Mn3+

(Table I).


Mn(111) -Organic Acids as Reductants

In solution, the gradual fading of the color complex indicates that the

organic acid is being oxidized to CO, and H 2 0 by the bound Mn( 111) while

Mn(111) is being reduced to Mn( 11). The rates of formation of the color

complex and its fading both are inversely proportional to the pH. When

redox decomposition such as this takes place in a soil, base-forming cations

are released, causing the pH to rise and the rate of oxidation of the organic

to be slowed or stopped. Mn(III)-citrate made from K,.,-citrate at pH 4.7

decomposes fairly rapidly for a few hours until the pH reaches 7.6, and

then it remains stable, with no more fading or C02 loss for several months.

It was pointed out in Section VI,C and Fig. 3 that the supermanganese

Mn3+ion can function as a double agent, powerful not only as an oxidizer

but also as a reducer. Mn(II1)-citrate will reduce methylene blue and

tetrazolium blue and oxidize tetramethylbenzidine (see Section XIV,D). In

solutions and in soil in the laboratory, under normal fluorescent lighting,

we found that Mn(II1)-citrate or oxalate, formed by reverse dismutation,

reduced Cr(V1) much more effectively at pH 4-6 than the organic acid



Figure 9. The marked lowering of net Cr oxidized, shown when Mn(I1) and oxalate ware

added together to soil samples containing different levels of oxidized Mn, demonstrates the

reducing power of Mn(ll1) formed by reverse dismutation.

alone. We also showed that Mn(II1)-citrate reduced Cr(V1) faster than the

organic alone.

These effects were strongly borne out in treated soil samples in which the

chromium net oxidation test (see Section XIV,F) was used to characterize

the oxidative minus the reductive powers of the samples (Bartlett, 1988).

Figure 9 shows that adding Mn2+to soil samples already containing MnO,

decreased the net oxidation of Cr by the soil. Part of the reason may be the

temporary increase in positive charge (Fig. 7) in repelling Cr( 111). But the

most probable reason is the reducing effect of Mn( III), which will form by

reverse dismutation (see Section IX,C,2) when Mn2+ is adsorbed onto

MnO,. Adding citrate alone had less effect than Mn2+alone on reduction

of Cr(VI), as shown by the net test. However, adding both Mn2+ and

citrate together markedly increased reduction of Cr( VI), as indicated by

the net Cr oxidation test. Reduction by Mn(II1) formed by reverse dismutation most surely is the explanation for the huge lowering of the Cr

oxidation net test with the two added together.



Manganese (111) also has the ability to oxidize an organic compound by

single electron steps to form a reducing organic free radical. For example, a

carboxyl free radical (R-COO.) has a strong disposition for giving up its

remaining odd electron to act as a reducing agent.

4. The Oxymoron: Enhanced Oxidation by Oxygen Restriction

In a slightly reducing environment, the first electron step in the partial

reduction of 0, is the formation of superoxide, the oxidizing/reducing free

radical. Manganese(III), formed in the first electron step in the partial

reduction of MnO, , is another extremely reactive oxidizing/reducing free

radical. Its reoxidation produces highly reactive “fresh” MnO, . This effect

is demonstrated by Mn behavior in oxidizing Cr in soils. For example,

partially restricting aeration by stoppering a flask escalated the Mn oxidizing behavior. Ten times the concentration of Cr(VI) was produced in a soil

suspension incubated 9 days with MnSO, and Cr(OH), in a stoppered

flask, as compared with the same volume of suspension swirled in an open

flask. Vigorous aeration of a high-organic-matter soil increased net Cr

oxidation by Mn oxides, but the same aeration lowered oxidation by a

low-organic-matter soil. Stoppering the low-organic-matter soil increased

net Cr oxidation, whereas stoppering the high-organic-matter soil halted

oxidation entirely. Thus, the redox poise between reactivities of oxidants

and reductants is critical. (R. J. Bartlett, unpublished data). Changing the

balance may reverse the direction of whatever is happening.

Compaction of soil in wheel tracks in turf typically results in dark green

stripes on the surface, not between areas of compaction, but at the site of

compaction. Chemical analysis of the vegetation shows that the extra green

is associated with higher nitrogen. The increased nitrogen is the result of

increased oxidation in the root zones of plants where the soil has been

somewhat compacted. Thus, it appears that partial exclusion of oxygen by

compacting a soil can increase certain oxidative processes in that soil.

The favorable oxidative effects resulting from restricting aeration seem

to be mainly related to the reactivity of Mn oxides. Fresh Mn oxides

provide better aeration than air in an oxidizing soil environment. The

oxides are more ready electron acceptors than oxygen. In paper towels

impregnated with high-Mn-oxide soil compared with those with only nutrient solution, white clover seeds germinated more quickly and had a

higher percentage of germination (R. J. Bartlett, unpublished data).

When aeration and respiration and synthesis of electron donors get out

of balance in the rhizosphere, dangerous-to-life free radicals may form.

The enzymes that scavenge such free radicals depend on metals, generally,

Cu, Mn, and Fe, to effect electron transfers. An example is a superoxide



dismutase described by Fridovich ( 1979, containing Cu2+and also Zn2+,

as a stabilizer. Reduced Mn acts as a scavenger for oxidizing oxygen free

radicals, and oxidized forms of Mn together with Fe are responsible for

accepting electrons from highly reactive reduced toxic or allelopathic organic compounds and then using these compounds for the synthesis of

benign, nurturing, and stable humic substances. Manganese does these

things in the dark, where plant roots grow and develop in soil. Roots

benefit from the humified materials surrounding them with the right balance between 0, and H,O,while, at the same time, they are being protected by redox metals from toxicity of oxygen free radicals.

5 . Mechanism for Oxidation of Manganese in Soils

Manganese(IV) usually occurs in the soil as a colloidal solid oxide. Often

the negative charges on the oxide surface are occupied by adsorbed Mn(II),

and this gives the overall surface a positive charge (Loganathan et al.,

I977), enabling its adsorption by negatively charged colloidal organic matter (Fig. 7). The change in surface charge of MnO, from negative to

positive is easily observed by a reversal in direction of electrophoretic

mobility (Bartlett, 1988). The positive charges also could arise from adsorbed Mn( 111) ions formed by reverse dismutation of Mn( 11) and MnO,.

It is axiomatic that living plants, animals, and microorganisms supply all

of the soil organic substances that are redox reactive. Soil microorganisms

are indispensable in synthesizing and making available phenolic and aliphatic acids and in influencing pH near reactive surfaces by mineralizing

organic matter and releasing base-forming cations. They also “graze” and

metabolize selectively the most biologically available electron-rich substances, those that would tend to interfere the most with oxidation of Mn.

Because the autooxidation of Mn(I1) by atmospheric 0, cannot be

demonstrated unless the pH is above 8 (Diem and Stumm, 1984), it has

been a common assumption that formation of Mn oxides in most soils

requires specific microbial enzymes, and activities of soil microorganisms

have been studied in this regard [e.g., Ehrlich ( 1 976), Silver et al. (1986),

and Sparrow and Uren (1987)l. Unfortunately, lack of Mn oxide formation after use of a chemical microbial inhibitor has been incorrectly used as

conclusive evidence for dismissing the importance of abiotic mechanisms

of Mn oxidation. Chemical inhibitors (e.g., chloroform or sodium azide)

will reduce MnO, in any soil containing organic acids and will destroy the

soil’s Mn-oxidizing mechanism (Ross and Bartlett, 198I). Using Cr oxidation to evaluate Mn oxides, Ross and Bartlett ( 1 98 I ) , showed that oxidation of added Mn was proportional to existing oxides. Arrhenius plots of

rates of Mn oxidation at different temperatures were indicative of nonbio-



logical characteristics for the oxidation. Ross and Bartlett (198 1) hypothesized that the oxidation was autocatalytic and that fresh Mn oxides tended

to form on old oxide surfaces.

Chemical oxidation of Mn(I1) added to an acid (pH 4.4)soil and to a

neutral soil was demonstrated directly in another study (Bartlett, 1988). All

of the oxidation of added MnSO, that was to occur during a 36-hr period

occurred the first 15 min. Obviously the oxidation was dependent on the

biochemical status of electron acceptors already present in the soils at the

moment of addition and not on microbial growth in response to the

Mn(I1) additions. After 5 days, there were marked increases in oxidized

Mn in both soils, suggesting that a biochemical readjustment had taken

place, presumably in response to the newly formed Mn(IV) and/or excess

added Mn(II), and to changes in microbial activities. It is not safe to

assume that Mn not extractable by a neutral salt has been oxidized because

strong inner sphere binding of Mn( 11)by soil organic matter above pH 5 - 6

can prevent its exchangeability (McBride, 1982).

Probably the hypothetical “manganese oxidase” enzyme of Silver ef al.

(1986) in reality is either the hydroxyl free radical (OH - ) or the protonated

superoxide free radical (HO;). These are the most likely electron acceptors

in the oxidation of Mn(I1) to Mn(1V). Fresh, newly formed Mn oxides

usually are found in soil regions where oxygen free radicals are being

formed, at redox interfaces, in rhizospheres, and in regions where atmospheric 0, is in somewhat short supply. Free radicals form, and Mn(I1) is

oxidized in scavenging them. Even if there are not microbes that have

specific roles as manganese oxidizers, microorganisms nevertheless are the

ultimate setters of the scene, and, of course, the oxygen free radicals can be

considered to be indirectly the result of microbial activity. A Mn( 11) ion, in

reducing OH and becoming oxidized in the process, is destroying it and

is preventing a microaerophile, busy setting the scene, from being poisoned

in its own juice.

Atmospheric 0, is the terminal electron acceptor when an oxygen free

radical oxidizes Mn( 11) to Mn( IV). When the soil pH and pe are both high,

as in the presence of free CaCO,, 0, may oxidize Mn(I1) to Mn(1V)

directly. Microbial processes also favor direct oxidation in high-pH microsites as they increase pH by releasing base-forming cations during decomposition of organic residues.

a ,



There are many bits of circumstantial evidence indicating that Mn is

involved in soil nitrogen redox transformations, but there is little under-



standing of the processes involved. Manganese oxides are thermodynamically capable of oxidizing NH, and N, to nitrate, but there is no hard

evidence that this happens. Circumstantial evidence consists of good

correlations between nitrate production and content of Mn oxides in

incubated soils. The good correlations may have resulted because soil Mn

oxides retarded or prevented denitrification by oxidizing nitrite back to

nitrate as fast as it formed, or else Mn oxides perhaps prevented denitrification by scavenging readily available organic reducing agents, oxygen free

radicals, or Fe( 11).

Both the oxidation of NH,OH to nitrate and of nitrite to nitrate by

synthetic amorphous MnO, are easily demonstrated in the test tube (Bartlett, 1981b, 1988). Bartlett ( I 98 1b) showed that the amount of nitrate

formed in soils from added nitrite at 0.5”C was directly related to the net

Cr oxidized by the standard oxidation test (see Section XIV,F), that is, to

the net oxidizing ability of the soil Mn oxides. Nitrate formation and

MnO, reduction were stoichiometrically related in the presence or absence

of atmospheric 0,. When MnO,/NO? ratios were high, reduction to

Mn( 111) was mainly observed; when low, Mn( 11) was the reduced product

accompanying nitrate formation.


The Cr cycle, Fig. 10, begins where it has ended, with Cr in its least

mobile form, Cr( III), precipitated or tightly bound by a variety of ligands,

such as hydroxyls, humates, and phosphates, or, in its most inert forms,

substituting for two atoms of Fe in the magnetite structure, as FeCr204,or

for small amounts of octohedral A1 in clay minerals.

Like Al, Cr( 111) can be mobilized by low-molecular-weight organic acids

such as citrate. The chelated Cr3+ may then interact with negatively

charged MnO, and become oxidized to Cr(VI), the HCrO; ion in the

diagram. Some of the citrate ligands are recycled. If there is a surplus of

citrate, the Mn2+formed when the Cr was oxidized may react with surplus

MnO, and reversely dismutate to two molecules of Mn(II1)-citrate, according to Eq. (36). Highly reducing Mn(II1)-organic, when it forms, will

temporarily interfere with further Cr oxidation.

The next step is “dechromification,” or the reduction of Cr( VI) by

carbon reduced by the sun’s energy through photosynthesis. An intermediate species, such as Fe2+or S2-, reduced by carbon, can serve as the direct

electron donor. Direct sunlight may hasten the process of Cr(V1) reduction. It is theoretically conceivable that dechromification, like denitrifica-

Tài liệu bạn tìm kiếm đã sẵn sàng tải về

VIII. Free Radicals in Redox Processes

Tải bản đầy đủ ngay(0 tr)