ARCHIVES

OF RIOCHEMISTRY

Vol. 276, No. 2, February

AND

BIOPHYSlCS

1, pp. 405-414,

1990

Characterization of Reactions Catalyzed by Manganese Peroxidase from Phanerochaete chrysosporium’ Michael

D. Aitken’

and Robert

L. Irvine

Department of Civil Engineering and Center for Bioengineering Ukuersity of Notre Dame, Notre Dame, Indiana 46556

and Pollution

Control,

Received July 31,1989, and in revised form October 6,1989

Manganese peroxidase (MnP) is one of two extracellular peroxidases believed to be involved in lignin biodegradation by the white-rot basidiomycete Phanerochaete chrysosporium. The enzyme oxidizes Mn(I1) to Mn(III), which accumulates in the presence of Mn(II1) stabilizing ligands. The Mn(II1) complex in turn can oxidize a variety of organic substrates. The stoichiometry of Mn(II1) complex formed per hydrogen peroxide consumed approaches 2:l as enzyme concentration increases at a fixed concentration of peroxide or as peroxide concentration decreases at a fixed enzyme concentration. Reduced stoichiometry below 2: 1 is shown to be due to Mn(II1) complex decomposition by hydrogen peroxide. Reaction of Mn(II1) with peroxide is catalyzed by Cu(II), which explains an apparent inhibition of MnP by Cu(I1). The net decomposition of hydrogen peroxide to form molecular oxygen also appears to be the only observable reaction in buffers that do not serve as Mn(II1) stabilizing ligands. The nonproductive decomposition of both Mn(II1) and peroxide is an important finding with implications for proposed in vitro uses of the enzyme and for its role in lignin degradation. Steady-state kinetics of Mn(II1) tartrate and Mn(II1) malate formation by the enzyme are also described in this paper, with results largely corroborating earlier findings by others. Based on a comparison of pH effects on the kinetics of enzymatic Mn(II1) tartrate and Mn(II1) malate formation, it appears that pH effects are not due to ionizations of the Mn(II1) complexing ligand. (c 1990 Academic Press, Inc.

Much scientific interest has been focused recently on the lignin-degrading capabilities of the white-rot basid1 This project was funded by Occidental Chemical Corp., Grand Island, NY. ’ To whom correspondence should be addressed. Present address: Department of Environmental Sciences and Engineering, School of Public Health, IJniversity of North Carolina, Chapel Hill, NC 27599. 7400. 0003.9861/90

$3.00

Copyright C 1990 by Academic Press, Inc. All rights of reproduction in any form reserved.

iomycete Phanerochaete chrysosporium. Two types of extracellular peroxidase have been discovered to be produced by the organism, and at least one of these enzymes is believed to be directly involved in lignin degradation (1,2). The lignin-degrading enzyme, referred to as ligninase (2) or lignin peroxidase (1)) has been studied extensively and its properties have been reviewed (3-7). Less is known about the role of the second type of extracellular peroxidase, which has been referred to as manganese peroxidase (MnP)” because of its requirement for divalent manganese in carrying out peroxide-dependent oxidations (8,9). Reactions catalyzed by MnP in the presence of manganese and hydrogen peroxide include oxidation of various dyes and phenols (s-12), decarboxylation and demeth(ox)ylation of aromatic substrates (8, 10,13), and oxidative cleavage of a phenolic lignin model dimer (14). The need for manganese in MnP reactions has been shown to result from the enzyme’s ability to catalyze the oxidation of Mn(I1) to Mn(II1) in the presence of certain ligands, including pyrophosphate (9), tartrate (12), and lactate (15). The trivalent manganese complexes formed enzymatically have been shown to oxidize several of the aromatic substrates of MnP (12, 15). Chemically generated Mn(II1) complexes such as Mn(II1) lactate (15) and Mn(II1) malonate (14) have been shown to oxidize organic substrates of MnP independent of the enzyme. Thus it appears that the role of manganese in organic compound oxidations by MnP is to serve as a one-electron transfer mediator. In more recent reports on MnP, Gold’s research group (16-18) has described spectral characteristics of the enzyme and its oxidized states. These workers have demonstrated the ability of MnP to form the peroxidase intermediates compound I and compound II (17, 18), and to form the species compound III in the presence of ex3 Abbreviations used: MnP, manganese peroxidase; azinobis(3-ethylbenzothiazoline-6-sulfonic acid).

ABTS,

2,2’-

405

406

AITKEN TABLE Characteristics

AND

I

of Mn Peroxidase

Preparations MnP activity

Preparation

Urn1 ’

Ml M2 M3

0.59 0.90 0.52

6.9 56 23

U (mg protein))’ 216 467 235

Note. Codes were assigned arbitrarily for purposes of distinguishing among different enzyme preparations. A407/A280 represents absorbance ratio of heme to total protein. Enzyme units are defined in the text.

by one (17). Compound I could be reduced compound II by adding Mn(II), ferrocyanide, or various phenolic substrates; however, only divalent manganese and ferrocyanide could reduce compound II to native enzyme. Divalent manganese therefore appears to be required for complete enzyme turnover in the catalytic cycle. Our work with MnP began as an effort to evaluate its ability to oxidize aromatic pollutants. In doing so, we needed to understand better the effects of various reaction conditions on observed reactions. From earlier work described above and from unpublished work in this laboratory, it appears that most organic substrates oxidized by MnP are oxidized by Mn(II1) in a postenzymatic reaction. To characterize the nature of direct enzymatic events, we focused on the effects of reaction conditions on Mn(I1) oxidation. Steady-state kinetics of Mn(II1) tartrate and Mn(II1) malate formation by MnP are described in this paper. Hydrogen peroxide was found to decompose the Mn(II1) tartrate complex in a manner that is catalyzed by Cu(I1) and with a corresponding formation of molecular oxygen. Peroxide-dependent decomposition reduces the net accumulation of Mn(III), affecting the observed stoichiometry and (in some cases) the kinetics of enzymatic Mn(II1) complex formation. The net decomposition of hydrogen peroxide by MnP and manganese in the absence of a Mn(II1) stabilizing ligand is reported as well. Based on the results of these studies, it appears that most observed reactions of MnP can be explained in terms of a competition between enzymatic Mn(II1) formation and nonenzymatic decompocess peroxide electron to

sition

of Mn(II1)

MATERIALS

AND

by hydrogen

peroxide.

METHODS

Enzyme Preparations General procedures for culturing Phanerochaete chrysosporium VKM F-1767 (ATCC 24725), enzyme production, and enzyme purification by ultrafiltration and ion exchange are described elsewhere (19). All enzymes were from carbon-limited cultures. Properties of the MnP preparations used are summarized in Table I. Preparation Ml was a combination of two early-eluting peaks from ion exchange, and

IRVINE

behaved kinetically as a single enzyme [good linearity of LineweaverBurk correlations for the range of manganese (15250 ).&I) and hydrogen peroxide (O-40 NM) concentrations tested]. Preparations M2 and M3 were from a single early-eluting peak, and were purified further with two more passes through the ion-exchange column, using salt gradient elution (starting buffer 5 mM Na succinate, pH 5.5; final buffer 5 mM Na succinate, 200 mM NaCl, pH 5.5). None of the preparations contained any detectable ligninase activity.

Enzyme Assay The MnP assay used in this study involved oxidation of guaiacol. The assay reaction mixture contained 100 mM Na tartrate, pH 5.0,lOO +M MnSO,, 100 PM guaiacol, enzyme solution, and 50 PM hydrogen peroxide. Reactions were started by adding hydrogen peroxide and were quantified by monitoring the initial rate of increase in absorbance at 465 nm (12). Since the nature of the reaction product is unknown, one unit of activity was defined as an initial increase in absorbance of 1.0 per minute. Using data provided by Paszczynski et al. (12), this definition of a unit of MnP activity corresponds to 0.114 unit based on an assay involving vanillylacetone oxidation. The guaiacol assay was reproducible (relative standard deviation 3.9% for five replicate assays) and was linear in enzyme concentration over a range of concentrations from 0.02 unit/ml to at least 0.3 unit/ml.

Chemical Synthesis of Mn(III)

Tartrate

Method 1. This method was patterned after the method of Mn(III) pyrophosphate production described by Kenten and Mann (20). MnS04 (10 mM) and MnOl (12 mM) were mixed in 250 ml of 100 mM Na tartrate, pH 4.3. The mixture was stirred overnight; then the residual MnO,, was filtered away using a 0.2.pm membrane filter. The resulting reddish-brown solution contained approximately 0.1 mM Mn(II1) tartrate, based on an estimated extinction coefficient as described below. Method 2. This method was similar to method 1, but 100 mM disodium tartrate (unbuffered), 5 mM MnS04, and 5.8 mM MnOl were used. The filtered solution was frozen and then partially thawed to yield a solution containing approximately 1.8 mM Mn(III) tartrate. Method 3. Several milligrams of Mn(II1) acetate powder (Aldrich Chemical Co.) were mixed in 5-10 ml of spectrophotometric-grade methanol. The resulting mixture contained a substantial amount of colloidal matter until several drops of double-distilled water were added. A translucent reddish-brown solution resulted, with a slight amount of dark precipitate that settled rapidly to the bottom of the vessel. This solution was stable at room temperature and was used to prepare working solutions of Mn(II1) tartrate directly. The working solutions of Mn(II1) tartrate consisted of Na tartrate buffer at a desired pH, MnS04 at 50 to 200 PM to stabilize the Mn(II1) complex, and a small volume of the methanol solution. Mn(II1) tartrate solutions (up to 60 PM) prepared in this manner were stable throughout each experimental period.

Absorbance Spectrum and Extinction Mn(III) Tartrate

Coefficient of

The ultraviolet absorbance spectrum of Mn(II1) tartrate was similar to that reported for Mn(II1) lactate (15), with peak absorbance at 236 nm. Due to some interference by Mn(I1) below 236 nm, absorbance of Mn(II1) tartrate was routinely monitored at 240 nm. An extinction coefficient at 240 nm was estimated by calibrating the concentration of the Mn(II1) acetate in methanol stock, then preparing a standard absorbance curve for various dilutions of the methanol stock in tartrate buffer (100 mM, pH 5.0). The concentration of the Mn(III) acetate stock solution was calibrated by titrating various volumes of the solution versus 200 KM potassium ferrocyanide in distilled water (ferrocyanide in excess). Reaction of Mn(II1) with ferrocyanide was as-

CHARACTERIZATION

OF MANGANESE

sumed to occur in a 1:l stoichiometry to produce Mn(I1) and ferricyanide. Ferricyanaide formation was followed at 420 nm, and a linear correlation was observed between ferricyanide formed and volume of Mn(II1) acetate stock added per milliliter of reaction mixture (r2 = 0.998, n = 6). The range of Mn(II1) concentrations tested, as well as the concentration of ferrocyanide used, was limited by precipitation of Mn(I1) ferrocyanide, which interfered with absorbance measurements. An extinction coefficient of 8.1 X 10s Mm’ cm-i at 240 nm was estimated from the linear standard absorbance curve (r’ = 0.9996, n = 7). This is probably an upper limit for the extinction coefficient because possible disproportionation of Mn(II1) was ignored in the ferrocyanide titration experiments [i.e., Mn(II1) concentration in the methanol stock may have been underestimated]. The extinction coefficient of Mn(II1) tartrate was observed to decrease as pH decreased, but this decrease was significant only below pH 4.2.

Stoichiometry

of Enzymatic

Mn(IIIJ

Tartrate

Formation

Enzymatic formation of Mn(II1) tartrate was verified by comparing ultraviolet absorbance spectra of the species formed enzymatically upon addition of hydrogen peroxide (in 100 mM Na tartrate, pH 5.5, containing 500 pM MnSOJ with that synthesized chemically by methods described above. Experiments to determine the stoichiometry of Mn(II1) formed versus peroxide consumed were run in reaction mixtures containing 100 mM Na tartrate, pH 5.0, plus enzyme, 250 PM MnSO,, and hydrogen peroxide. Enzyme and peroxide concentrations used are provided in the text. Reactions were run until a stable value of A240 was reached. Preliminary experiments indicated that conversion of virtually all of the oxidizing equivalenm of hydrogen peroxide initially added to the reaction mixture were accounted for in the final products [Mn(III) complex and molecular oxygen].

Enzyme

Kinetics

Experiments

Kinetic parameters for enzymatic formation of Mn(II1) tartrate and Mn(II1) malate were determined from initial velocity experiments using Lineweaver-Burk correlations. Mn(II1) tartrate formation was followed at 240 nm and Mn(II1) malate was followed at 245 nm (wavelength of peak absorbance). Reaction mixtures contained 100 mM buffer (tartrate or malate) at the desired pH, plus enzyme, MnSO,, and hydrogen peroxide at concentrations given in the text. Reactions were started by adding peroxide. Concentrations of peroxide used in kinetic experiments were somewhat lower than optimal; normally, a range of substrate concentrations above and below the Km should be evaluated in enzyme kinetics studies. Due to the inhibitory effects of high peroxide concentration on MnP activity (9, 12; discussed further below), peroxide concentrations were limited to a range near the K,,, and below. Satisfactory results (good linearity of LineweaverBurk correlations) were obtained in all cases,

Spectrophotometry Optical absorbance spectra and kinetic measurements were made using a Varian DMS 100 scanning spectrophotometer (Varian Instrument Division, Sunnyvale, CA). The instrument was converted to a Model DMS 200 during the latter part of this project, which allowed automatic determination of the reaction rate (dA/‘dt) at any time during the course of a reaction. Kinetics of Mn(II1) tartrate formation (enzymatic systems) or decomposition (chemical systems) were measured by following initial rates of change in absorbance at 240 nm. Ox.~grn euolution. Oxygen production from hydrogen peroxide during certain MnP reactions was followed in a stirred, water-jacketed cell (Gilson Medical Electronics, Middleton, WI) fitted with a Clarktype oxygen electrode. Water at a constant temperature of 23°C was circulated through the cell with a constant-temperature circulator. The oxygen electrode was connected to a digital amplifier/meter (Yellow Springs Instruments, Yellow Springs, OH) and was calibrated us-

PEROXIDASE

407

REACTIONS

ing air-saturated double-distilled water immersed in the water bath of the circulator. An oxygen saturation value at 23°C of 272 PM (21) was used for calibration. Buffers and reagents were placed in the sample cell and allowed to equilibrate with stirring to a steady dissolved oxygen concentration before reactions were started (by adding enzyme or hydrogen peroxide, as appropriate). Initial rates of oxygen production were estimated by drawing a straight line through recorded traces of dissolved oxygen concentration versus time. The same apparatus was used to measure residual hydrogen peroxide concentrations when desired by introducing catalase to the reaction system.

Reagents Buffers used in studies to determine kinetic parameters as a function of pH were prepared by mixing 100 mM solutions of organic acid with 100 mM solutions of the disodium salt of the organic acid. The pH meter was calibrated using commercial reference standards bracketing the pH range of interest. Accuracy of the pH of Na tartrate buffers was 0.01 pH unit. Accuracy of Na malate buffers was estimated to be within 0.05 pH unit. Catalase (bovine liver) was purchased from Sigma Chemical Company (St. Louis, MO) and Mn(II1) acetate was purchased from Aldrich Chemical Company (Milwaukee, WI). All other reagents were ACS reagent grade or purer. Stock solutions of hydrogen peroxide were prepared from 30% reagent, calibrated by titration against KMn04 stanbottle at 4°C. Hydard in 1 N H,SO,, and stored in a paper-wrapped drogen peroxide stock solutions were stable for several months, as judged by periodic recalibration. Working solutions of peroxide were prepared by dilution of stock on the day of use. Reagent water used to prepare and dilute reagents was double distilled.

Data Analyses Curve fitting was conducted by linear or nonlinear regression analyses (least-squares minimization) on a personal computer using the equation-solving program TK! Solver Plus (Universal Technical Systems, Rockford, IL). Simultaneous differential equations were solved with the same software using a fourth-order Runge-Kutta technique. RESULTS

Steady-State

Kinetics

of Mn(III)

Tartrate

Formation

Steady-state kinetics of enzymatic Mn(II1) tartrate formation were evaluated as a function of hydrogen peroxide concentration and divalent manganese concentration. Double-reciprocal plots are shown in Fig. 1. From the data presented in Fig. la, a K,,, value for hydrogen peroxide at saturating manganese concentration (250 ,uM) was estimated to be 59 PM for enzyme preparation M2. A similar K,,, for hydrogen peroxide (57 FM) was observed for enzyme preparation Ml, which was isolated from a production batch different than that used to isolate preparation M2. These K, values for peroxide are significantly lower than the value of 140 PM reported by Glenn et al. (15) for formation of Mn(II1) lactate by the enzyme. The lines of best fit in the double-reciprocal plots of Fig. la are parallel (standard deviation of the slopes 4.5% of the mean). Replots of data from Fig. la with hydrogen peroxide as fixed substrate are shown in Fig. lb. Reciprocal velocity at fixed peroxide concentration is plotted against the reciprocal of manganese concentration squared. The

408

AITKEN

AND

IRVINE

lation for preparation of 7.1 X lo3 min-l. Effect of pH on Mn(IIIJ

30 .A l 7

=;‘ 20 ? s % - 10

0’ 0.0

I

I

/ 0.2

0.1

[HzOzl-l, PM-’

01 0

I_ I 0.001

I 0.002 [Mn]-2,

I 0.003

I 0.004

I 0.005

Ml resulted in a turnover

Complex Formation

number

Kinetics

Using fixed reaction conditions (single rate observations), Glenn et al. (15) found that Mn(II1) lactate formation by MnP reached a maximum rate at pH 5.0, and that ABTS oxidation rate had a maximum at pH 4.5 (no data were reported between pH 4.5 and 5.0). Huynh and Crawford (10) observed peak oxidation of vanillylacetone in tartrate buffer at pH 5.0. These earlier findings were extended in this study by determining the effect of for both Mn(II1) tartrate PH on Km,app03202) and Vmax,app and Mn(II1) malate formation. Plots of K,,a,,(H,O,) versus pH are shown in Fig. 2. It appears that two ionizable groups affect the observed K,,, values in both the tartrate and malate systems, and in both cases the maximum Km,appwas observed to be at about pH 4.8. Similar patterns were observed for V max,app values in both buffer systems. Based on the similarity of pH effects in the two buffers, it appears that such effects are not a result of ionizations of the Mn(II1) complexing ligand. Other possible explanations for pH effects include ionizations at the enzyme active site (23) or dissociation of heme from the apoprotein (24). It should be noted that further effects of pH may be observed when organic compounds are oxidized via Mn(III), because of pH effects on the rate of reaction between Mn(II1) and organic substrates [unpublished observations; also see Ref. (25)]. The parameter Km,app(HSOa)/Vmax,appwas essentially independent of pH in both buffer systems (standard de-

pK2

FIG. 1. Lineweaver-Burk plots of Mn(II1) tartrate formation by MnP. Preparation M2 (207 mU ml-‘) was used in 100 mM Na tartrate, pH 5.0. (a) Plots with Mn(I1) as fixed substrate at concentrations of 15 (0), 20 (0), 30 (Cl), and 250 (a) PM. (b) Replots with peroxide as fixed substrate at concentrations of 5.1 (O), 6.6 (O), and 7.8 (a) FM, and V& (0) (from plot a) versus [Mn(II)j-‘.

correlation of these plots was significantly better than plots of reciprocal velocity versus reciprocal manganese concentration. A K,,, value of 36 FM for manganese was determined from Vi’,, versus [Mn(II)le2. Using the V,,, value obtained at saturating manganese concentration from Fig. la, the extinction coeffcient for Mn(II1) tartrate, and a molecular weight for MnP of approximately 46,000 (9, 12, 22), the minimum turnover number for Mn(II1) tartrate formation was estimated to be 5.9 X lo3 mol Mn(II1) tartrate formed per mole enzyme per minute (this assumes all protein in preparation M2 was MnP, which probably was not the case based on its ratio of A407/A280). An equivalent calcu-

80.

I

1

0

I

0

0

4.0

I 5.0

I

L 6.0

PH FIG. 2. Effect of pH on K, (HZ02) for Mn(II1) tartrate (0) and Mn(II1) malate (a) formation by MnP. Preparation M2 (107 mU ml ‘) was used in 100 mM solutions of the complexing ligand containing 250 FM MnS04.

CHARACTERIZATION

a 4 z T;: 2 5 g

-2 ; g T E

OF MANGANESE

PEROXIDASE

409

REACTIONS

ide concentration were used, Mn(II1) tartrate would reach a maximum concentration and then would decrease with further reaction time. Similar decreases in Mn(II1) tartrate concentration with time were observed if a high concentration of peroxide was added to a mixture containing Mn(II1) tartrate previously formed enzymatically.

2.2 2.0 1.8

1.6

Decomposition 1.4 1.2 1.0

10

20

[H,O,]

30

40

added, PM

FIG. 3. Stoichiometry of Mn(II1) tartrate produced as a function of initial concentration of hydrogen peroxide added. Enzyme preparation Ml at 35 mlJ ml ’ (0) or 207 mU ml-’ (a) was used in 100 mM Na tartrate, pH 5.0, containing 250 pM MnSO,. The solid and dashed lines represent numerical solutions to Eqs. [Z] and [3] using experimentally determined values of K, (H,O,), V,,, (for the lower and higher enzyme concentration, respectively), and kinetic coefficients k and n for Mn(II1) tartrate decomposition by peroxide.

viation 6.3% of mean for tartrate from pH 4.0 to 5.8, and 9.2% of mean for malate from pH 4.2 to 6.0). This result indicates that reaction of hydrogen peroxide with native enzyme is not a rate-limiting step (26) in manganese oxidation by MnP, and is consistent with observations made on other peroxidases (27,28). It is also consistent with observations that the rate of compound I formation in MnP is independent of pH (18). Stoichiometry

of Mn(III)

Tartrate Formation

The stoichiometric ratio of Mn(II1) tartrate produced per hydrogen peroxide added is shown in Fig. 3. As shown, the stoichiometric ratio varies with both the enzyme concentration and the initial peroxide concentration. At very low peroxide concentration, Mn(II1) produced per peroxide consumed approaches 2.0. As peroxide concentration increases, however, the stoichiometric ratio decreases. These decreases are not as great at higher enzyme concentration. From independent experiments (29), enzyme inactivation would not have been extensive over the range of conditions used in stoichiometry tests. Final Mn(II1) formation stoichiometry also was reduced as tartrate concentration was reduced below 50 mM under otherwise typical reaction conditions. The system used to prepare the data in Fig. 3 contained an initial concentration of Mn(I1) in stoichiometric excess for all peroxide concentrations tested. If initial Mn(I1) concentrations less than twice the initial perox-

of Mn(III)

Tartrate

The most likely explanation for the phenomena observed in enzymatic reaction mixtures described above is that Mn(II1) tartrate can be decomposed by hydrogen peroxide (30). Such an explanation was confirmed by observing the production of molecular oxygen under conditions in which Mn(II1) decomposition occurred. Additional loss of Mn(II1) tartrate over time might also result from spontaneous decomposition. Both spontaneous decomposition of Mn(II1) tartrate and decomposition by hydrogen peroxide were evaluated in enzyme-free systems. Spontaneous decay was observed to be first order in Mn(II1) tartrate concentration over a pH range from 3.0 to 5.0, as reported in Table II. The reaction systems used to evaluate spontaneous decay of Mn(II1) tartrate contained some Mn(I1) [from the Mn(II1) tartrate preparation], so that it is not known to what extent the decay rates reported may have been affected by Mn(I1); divalent manganese is known to stabilize Mn(II1) in aqueous solution (31). Decomposition of Mn(II1) tartrate by hydrogen peroxide was much more significant than was spontaneous decay over the pH range of interest. The initial rate of Mn(II1) tartrate decomposition as a function of HzOz and Mn(II1) concentrations at pH 4.75 is shown in Fig. 4. Analysis of the data from Fig. 4 indicated a first-order reaction with respect to peroxide but a reaction order between 1 and 2 with respect to Mn(II1) tartrate concentration. The actual order was determined by assuming the following kinetic expression: TABLE

II

Spontaneous Decay Rate of Mn(II1) Tartrate as a Function of pH

PH 3.0 4.0 5.0

First-order constant 0.49 0.17 0.066

decay (h- ‘)

r2 (n) 0.988 (6) 0.993 (6) 0.999 (5)

Note. Mn(II1) tartrate was prepared in concentrated form by method 2 (see Materials and Methods). Aliquots of the concentrated solution were added to 100 mM Na tartrate at indicatedpH. Initial rate of decay was followed spectrophotometrically. Correlation coefficients were computed from plots of initial rate vs initial Mn(II1) concentration.

410

AITKEN

[H,O,] added,

AND

IRVINE

linked and can be modeled by a pair of differential equations:

PM

d[Mn(III)]/dt d [H202]/dt

-15f FIG. 4. Rate ofdecomposit.ion of Mn(III) tartrate by hydrogen peroxide as a function of peroxide concentration and Mn(II1) tartrate concentration. Mn(II1) tartrate was prepared by method 3 (see Materials and Methods) in 100 mM Na tartrate, pH 4.75, containing 50 GM MnSO,. Initial Mn(II1) tartrate concentration was 15.3 (0), 29.0 (O), or 55.6 (A) pM.

d[Mn(III)]/dt

= -k[Mn(III)]“[H,O,]

111

where k[Mn(III)]” = observed first-order reaction rate coefficient at constant [Mn(TII)]. A log-log correlation of the observed first-order rate coefficient versus Mn(II1) tartrate concentration (r2 = 0.997) resulted in a value for n of 1.33 and a value for k of 8.35 X 10’ M-““3 s i. Data shown in Fig. 4 were from systems in which 50 PM divalent manganese was present. Earlier studies with Mn(II1) tartrate in the presence of millimolar concentrations of Mn(II) indicated that t,he rate of decomposition was approximately first order in Mn( III) concentration. This result further illustrates the stabilizing effect of Mn(I1) on Mn(II1) decomposition. The rate of Mn(II1) decomposition by hydrogen peroxide increases with decreasing pH. Apparent secondorder rate constants for Mn(II1) tartrate decomposition [initial rate divided b y initial Mn(II1) and peroxide concentrations] determined at 30 PM Mn(II1) were 69.9, 49.6, 27.0, and 22.6 M-’ 5-i at, respectively, pH 3.5, 4.0, 4.75, and 5.25. Except for pH, reaction conditions were the same as described in Fig. 4. Effect of Mn(III) Stoichiometry

Decomposition and Kinetics

= uenz- (k[H,Oz][Mn(III)]“)

= -O.~(V,,,) - 0.5(k [H,Oz] [Mn(III)]“)

[2] [3]

where u,,, is the enzymatic rate of Mn(II1) formation modeled by a Michaelis-Menten equation for saturating Mn(I1) (single-substrate equation), and the parameters k and rz are as described above. These simultaneous differential equations were solved numerically to determine the final stoichiometry of Mn(III) produced per hydrogen peroxide consumed as a function of initial peroxide concentration and enzyme concentration. Results from the numerical analysis are shown in Fig. 3 along with the experimentally determined stoichiometric ratios. The possibility that Mn(III) complex decomposition may have affected initial velocity measurements with MnP was evaluated. Using an analysis similar to that described above for the effect of Mn(II1) decomposition on stoichiometry, the effect of Mn(II1) decomposition on initial velocity studies was determined to be negligible over the range of conditions used in the experiments reported in Figs. 1 and 2. Inhibition

of MnP by Cu(II)

Various metal ions, including Co(II), Cu(II), Fe(II), and Fe(W), have been reported to inhibit manganesedependent oxidations of organic compounds by MnP (12). The effect of metal ion inhibitors on formation of Mn(II1) complexes by MnP has not been reported.

[H,O,]

0

10

added,

/.LM

20

on Observed

Consumption of hydrogen peroxide in the enzyme system is related to both enzymatic Mn(II1) formation and Mn(II1) decomposition. Thus, overall rates of Mn(II1) formation and hydrogen peroxide consumption are

FIG. 5. Efkect of Cu(I1) on rate of Mn(II1) tartrate decomposition by hydrogen peroxide. Mn(II1) tartrate (prepared by method 3) concentration was 21 LLM in 100 mM Na tartrate, pH 4.75, containing 50 pM MnSO, and C&O4 at 0 (0) or 100 (A) FM.

CHARACTERIZATION

OF MANGANESE

r

0’

I 0.02

0

I

0.04

[H&1-‘,

I

I

0.06

/X1

FIG. 6. Eflect of Cu(I1) on kinetics of enzymatic Mn(II1) tartrate formation by MnP. Preparation Ml (104 mU ml’) was used in 100 mM Na tartrate, pH 5.0, containing 20 FM MnS04 and CuSO, at 0 (0) or 50 (A) pM. Corrected data accounting for Mn(III) decomposition by peroxide in the presence of copper are also shown (X).

In preliminary experiments, Cu(I1) was observed to inhibit both the initial rate and the final stoichiometry of enzymatic Mn(II1) tartrate formation. Oxygen production also was observed during the course of enzymatic oxidation of Mn(I1) in the presence of copper. In further testing, Cu(I1) was found to stimulate the nonenzymatic reaction of Mn(II1) tartrate with hydrogen peroxide. Involvement of copper in the nonenzymatic reaction was verified by measuring the rate of peroxidedependent decomposition of chemically generated Mn(II1) tartrate in both the presence and the absence of Cu(I1). Results are shown in Fig. 5. The observed second-order decomposition rate coefficient [slopes of the lines in Fig. 5 divided by the initial Mn(II1) concentration] in the presence of copper was an order of magnitude higher than in its absence (5.8 X 10’ Me’ s-l versus 87 M- ’ s-l). Initial velocity studies on enzymatic formation of Mn(II1) tartrate were affected when copper was added to the reaction system. Results from one experiment are shown in Fig. 6. Lineweaver-Burk plots are roughly parallel, but this resulted from reduced observed initial reaction velocities in the presence of copper. Observed initial rates in the presence of Cu(I1) were corrected for Mn(II1) decomposition by estimating the loss of Mn(II1) over the observation period used for each data point. Corrected data are plotted in the same figure, and match quite well the data obtained in the absence of copper. Mokcular

Oxygen

Evolution

in Alternative

PEROXIDASE

411

REACTIONS

presence of a Mn(II1) stabilizing ligand. Glenn and Gold (9) illustrated the effect of different buffers on the ability of MnP to catalyze the manganese-dependent oxidation of ABTS. Among the buffers in which ABTS oxidation occurred, malate (30, this work), tartrate (this work), citrate (30), and lactate (15, 30) have all been demonstrated to form relatively stable Mn(II1) complexes. Buffers in which ABTS oxidation did not occur included acetate, 3-hydroxybutyrate, pyruvate, and succinate (9). Glenn and Gold (9) also demonstrated that pyrophosphate inhibits ABTS oxidation in lactate buffer. This inhibition is in accord with the highly stabilizing effect of pyrophosphate on Mn(II1) (31, 32), presumably to such an extent that Mn(II1) pyrophosphate reacts more slowly with ABTS than does Mn(II1) lactate. Monitoring the production of molecular oxygen was a useful probe for systems in which stable Mn(II1) complexes could not be observed, nor in which manganesedependent oxidation of organic substrates occur. Oxygen production in these systems was indicative that some kind of enzyme activity was still occurring. Oxygen was not produced in any system under any of the experimental conditions tested in which manganese was omitted from the reaction system. In addition to experiments using tartrate as a buffer and complexing ligand for Mn(III), succinate, formate, and acetate were tested for their effects on oxygen evolution in enzyme systems containing divalent manganese and hydrogen peroxide (at pH 4.75). Oxygen production was observed in succinate buffer under a variety of conditions in which Mn(II), Cu(II), and buffer concentrations were varied. In all cases, the final concentration of oxygen produced approached 50% of the initial concentration of hydrogen peroxide added to the reaction system. However, the initial rate of oxygen production varied depending on the reaction conditions, as shown in Table III. At low (10 mM) succinate concentration, the

TABLE

III

Effect of Reaction Conditions on Oxygen Production Rate in MnP/Succinate Systems Succinate concentration (mM) 10

10 100 100 100

100 100

lMn(II)l

[CuUI)I

(PM)

(N)

60 60 60 60 60 500 500

0 100 0 100 300 0 100

Initial O2 production rate (gM min -‘)

0.6 1.2 2.3, 1.9 0.4 1.0 2.0 2.0

Buffers

The ability of manganese to serve as a mediator in organic compound oxidations by MnP may depend on the

Note. Reactions were carried out at pH 4.75 using 170 mU ml ’ enzyme preparation M3. Reactions were started by adding 29 WM hydrogen peroxide. Duplicate values are from duplicate tests.

412

AITKEN

AND IRVINE

presence of Cu(I1) increased the initial rate of oxygen evolution. At 100 mM succinate and 60 PM Mn(II), the initial rate of oxygen production was substantially higher in the absence of Cu(I1) than in the presence of 100 PM Cu(I1). This effect of copper was not observed at saturating Mn(I1) (500 PM), in which the initial rate of oxygen evolution was equally high in the presence and absence of copper. Also, in the absence of copper, the initial rate of oxygen evolution was higher at the higher succinate concentration. In a 50 mM succinate:50 mM tartrate mixture (pH 4.75), oxygen production rate in the absence of copper was low, comparable to that obtained in 50 mM tartrate alone. Addition of Cu(I1) to the 50:50 buffer mixture system resulted in a high rate of oxygen production, again comparable to that obtained in 50 mM tartrate alone (but contrary to the effect of succinate alone). Thus it appears that succinate does not participate significantly in the oxygen production reaction(s) when tartrate is present in the system. Oxygen was not produced at all in formate buffer under any reaction conditions. Thus, formate was considered to inhibit manganese-dependent reactions leading to oxygen evolution. Residual peroxide concentrations in these systems were not determined because formate inhibited catalase as well (even though catalase worked quite well in other buffers under similar conditions). Acetate was similar to formate in its effect on oxygen production, except at low concentrations (less than 10 mM). At 100 mM acetate, no response was observed under any reaction conditions. With acetate, unlike formate, it was possible to demonstrate that lack of oxygen evolution corresponded to lack of peroxide consumption; addition of catalase to reaction mixtures after a 15-min reaction period resulted in virtually complete recovery of the initial amount of hydrogen peroxide added. At acetate concentrations of 10 mM and below, oxygen production rates were low, but measurable, and increased as acetate concentration decreased. Copper(U) increased oxygen production rates significantly, but the effect was more pronounced at high Mn(I1) concentration, as shown in Table IV. DISCUSSION

The parallel double-reciprocal plots shown for divalent manganese as fixed substrate in Fig. la are typical of “ping-pang” kinetics for two-substrate enzyme systems (26). Such a result is consistent with the classical peroxidase mechanism (33), which in its simplest form for MnP can be written as (18) MnP + HzOz + compound I compound I + Mn(I1) compound II + Mn(I1)

[41

+ compound II + Mn(II1)

[5]

+ MnP + Mn(II1).

[61

TABLE IV

Effect of MnUI) and Cu(I1) Concentrations on Oxygen Evolution in MnP/Acetate Systems [Mn(II)I

[Cu(II)I

(PM)

(PM)

Initial O2 evolution rate (FM min-‘)

60 60 500 500

0 100 0 100

0 0.65,0.76 0.34,0.34 2.55,2.34

Note. React,ions were carried out in 5 mM Na acetate, pH 4.75, containing 170 mU ml-’ enzyme preparation M3. Reactions were started by adding 29 pM peroxide. Duplicate values shown are from duplicate tests.

Manganese-dependent turnover of MnP has been observed in succinate buffer (16-18), in which a stable Mn(II1) complex does not seem to form. It therefore appears that stabilization of Mn(II1) is not a requirement for complete turnover of the enzyme in the presence of manganese. However, turnover of MnP compound I has been observed to be more rapid in the presence of a Mn(II1) stabilizing buffer (lactate) than in succinate (18). This has been attributed to slower rates of manganese dissociation from the enzyme in succinate buffer (18). The minimal mechanism of reactions [4]-[6] can be shown to result in parallel double-reciprocal plots when either substrate is the fixed variable. However, as indicated in Fig. lb, kinetic results with hydrogen peroxide as fixed substrate are best fit by plotting reciprocal reaction rate versus reciprocal manganese concentration squared. This is an unexpected result that is not consistent with the simple mechanism shown above. Such a result would be obtained, for example, if two Mn(I1) were bound to the enzyme consecutively before formation of Mn(III), or if the second Mn(I1) were bound before release of the first Mn(II1). The results should be considered empirical at present, although they were useful for estimating a K, for Mn(I1). The stoichiometry of Mn(II1) tartrate produced per hydrogen peroxide consumed increases as enzyme concentration increases for a given concentration of peroxide, approaching a value of 2:l. Glenn et al. (15) previously reported a stoichiometric ratio of 1.48 for formation of Mn(II1) lactate by MnP and a ratio of 1.80 for Mn(II1) pyrophosphate. A 2:l stoichiometry is expected because hydrogen peroxide is a two-electron oxidant. Stoichiometric ratios below 2:l can be explained in terms of Mn(II1) decomposition by hydrogen peroxide, as illustrated clearly in Fig. 4 and as previously reported by Archibald and Fridovich (30). Consequently, in the enzymatic system the reaction in which the Mn(II1) complex is formed enzymatically (while peroxide is be-

CHARACTERIZATION

OF MANGANESE

ing consumed) competes with the reaction in which the Mn(II1) complex is decomposed by peroxide. The extent to which Mn(II1) will accumulate in a batch reaction, and therefore the final stoichiometric ratio of Mn(II1) produced to peroxide consumed, depends on the relative rates of the two competing reactions. Since peroxide is consumed in both the formation and the decomposition of the Mn(II1) complex, the net reaction can be driven to accumulate more Mn(II1) by increasing its rate of formation, that is, by increasing the enzyme concentration, This phenomenon is illustrated in Fig. 4, which shows lines fit to Eqs. [2] and [3] using two different enzyme concentrations. In addition, the Mn(II1) complex will accumulate to a greater extent with more stable complexes (i.e., complexes that react more slowly with hydrogen peroxide). Reaction between Mn(II1) and hydrogen peroxide has been proposed to occur by the following mechanism (30): Mn”’ MnOl

+ HzOz = MnOl

+ 2 H+

[71

+ Mn3+ = (Mn~O-O-Mn)4f

@I

(Mn-O-O-Mn)4+

+ 2 Mn2+ + O2

2Mn3’ + H202 + 2Mn2+ + 2H+ + 0,.

PI [lOI

In the above mechanism, the intermediate MnO: represents a divalent manganese complex with superoxide anion (the corresponding protonated form is MnOOH2+) (31, 34, 35). Reaction [7] is shown as an equilibrium reaction because the reverse reaction is known to occur in the presence of ligands that stabilize Mn(II1) (31, 36). If the net forward rate of reaction [7] were a limiting step in the above mechanism, then the overall reaction would be first order in Mn(II1) concentration. Observed reaction orders were between 1 and 2 at pH 4.75, which may indicate that, at that pH, reaction [8] or [9] was at least partially rate limiting. Several metal ions have been observed to inhibit organic compound oxidations by MnP in the presence of divalent manganese (12). Of these metal ions, Cu(I1) was shown in this study to effectively inhibit MnP by catalyzing the decomposition of Mn(II1) tartrate in the presence of hydrogen peroxide. It is possible that other metal ions act to inhibit apparent MnP activity in the same way. The ability of Cu(I1) to catalyze the peroxide-dependent reduction of Mn(II1) can be understood by considering the mechanism for reaction of Mn(II1) and hydrogen peroxide shown above. It is possible that Cu(I1) reacts with MnOi as follows: Mn02f + Cu2+ + Mn2+ + Cu+ + 0,.

[Ill

Such a reaction may be considered to be analogous to reaction of Cu(I1) with superoxide (37). If reaction [8]

PEROXIDASE

REACTIONS

413

were a rate-limiting step in the decomposition of Mn(III), then an increase in the rate of reaction of MnOz (by reaction [II], for example) would be expected to lead to an increase in the observed Mn(II1) decomposition reaction rate. Beyond a certain concentration of copper, however, reaction [7] would become rate limiting, and the overall rate of decomposition of Mn(II1) would not increase with increasing Cu(II) concentration. Apparent saturating concentrations of Cu(I1) were observed for inhibition of enzymatic Mn(II1) tartrate formation (data not shown). Decomposition of Mn(II1) by hydrogen peroxide also explains other experimental observations. Initial rates of enzymatic Mn(II1) tartrate formation decreased as tartrate concentration decreased and corresponded to increasing rates of oxygen evolution (via reaction [lo]; data not shown). This result seems to indicate that reaction of Mn(II1) with hydrogen peroxide competes with the Mn(II1) tartrate formation reaction. In buffers that do not serve as stabilizing ligands for Mn(III), such as succinate and acetate, oxygen evolution via peroxide decomposition was the only observable reaction. In succinate, oxygen evolution rate increased as succinate concentration increased, which may indicate that succinate participates in the reaction leading to decomposition of Mn(II1) (decomposition of hydrogen peroxide). Also, in succinate the initial rate of O2 production did not appear to depend on starting Mn(I1) concentration at the higher succinate concentration tested (100 mM). The effect of Cu(I1) on initial rate of oxygen production in succinate was a complicated function of Mn(II), Cu(II), and succinate concentrations; however, Cu(I1) did not catalyze the net decomposition of hydrogen peroxide beyond the maximum rate observed in the absence of copper. This result conflicts with results from tartrate systems. Oxygen evolution was not observed in 100 mM acetate, and was observed in 5 mM acetate only at a high Mn(II) concentration (500 pM) or in the presence of Cu(I1). Reactions in acetate may be analogous to reactions in tartrate [straightforward competition between Mn(II1) formation and decomposition], except that acetate does not stabilize Mn(II1) and therefore Mn(II1) would not be expected to accumulate significantly. At acetate concentrations above 5 mM, the lack of any observable reaction may have been due to the nature of acetate complexation to divalent manganese. Acetate is known to be a bridging ligand for Mn(I1) (38), and such a coordination of Mn(I1) may have inhibited its binding or reaction at the enzyme active site. The net decomposition of hydrogen peroxide to form molecular oxygen in buffers that do not serve as Mn(II1) stabilizing ligands raises important questions as to the significance of the reaction environment for MnP in lignin biodegradation; if Mn(II1) complexes are important in lignin degradation, then a Mn(II1) stabilizing ligand

414

AITKEN

AND

must be made available during periods in which MnP is active. Such stabilizing ligands could include organic acids produced during sugar oxidation by the fungus. With respect to potential waste treatment applications of the enzyme, the use of a system requiring a high concentration of a Mn(II1) stabilizing ligand is not expected to be practical. Relevant substrates that allow for complete turnover of MnP in the absence of manganese must be investigated instead.

We thank Dr. Tom Nowak of the Department of Chemistry for a critical review of the manuscript, and Dr. Larry Patterson of the Notre Dame Radiation Laboratory for useful discussions on manganese chemistry.

2. Tien, M., and Kirk, T. K. (1983) Science 221,661-663. 3. Tien, M., and Kirk, T. K. (1984) Proc. N&l. Acad. Sci. USA 81,

2280-2284. 4. Renganathan,

V., Miki, K., and Gold, M. H. (1986) Arch. B&hem. Biophys. 246,155-161.

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H. E. (1987) Phil.

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K. E., Tardone, P. ,J., Moen, M. A., and Price, I,. A. (1989) Arch. Biochem. Biophys. 270,404&409.

26. Siegel, I. H. (1975) Enzyme Kinetics, Wiley, New York. 27. Dunford, H. B., and Stillman, J. S. (1976) Coord. Chem. Reu. 19, 187-251. 28. Tien, M., Kirk, T. K., Bull, C., and Fee, J. A. (1986) J. Biol. Chem. 261,1687-1693. 29. Aitken, M. D., and Irvine, R. L. (1989) Biotechnol. Bioeng., in press.

30. Archibald, F. S., and Fridovich, I. (1982) Arch. Biochem. Biophys. 214,452%463. 31. Cabelli, D. E., and Bielski, B. H. J. (1984) J. Phys. Chem. 88, 3111-3115.

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