Ligand-dependent redox chemistry of Glycera dibranchiata hemoglobin

362

Biochimica et Biophysica Acta 828 (1985) 362-368

Elsevier BBA32183

Ligand-dependentredox chemistry of Glycera dibranchiata hemoglobin A.W. Addison and S. Burman Chemistry Department, Drexel University, Philadelphia, PA 19104 (U.S.A.)

(ReceivedDecember6th, 1983) (Revised manuscript receivedDecember21st, 1984)

Key words: Hemoglobin;Redox potential; Histidine;(G. dibranchiata, annelid)

The trans-l,2-diaminocydohexane-N,N,N',N'-tetraacetatoferrate(lll)/(ll) couple was used as a redox buffer-mediator to determine that the redox potential of the monomeric hemoglobin of Glycera dibranchiata is +0.153 ±0.008 V at 27°C in 0.1 M (pH 7.0) phosphate. The value is similar to that for Aplysia myoglobin, which also lacks distal histidine. Correlation of E °' with log p t/2(02) is discussed. Hexacyanoferrate(II) binds to the protein, resulting in an unusual enhancement of oxygen affinity. The binding of azide anion to the aquomet-protein has been studied by a spectropotentiostatic method and by the conventional spectrophotometric technique, and the results are interpreted in terms of an ionic dependence. Introduction The low-molecular-weight hemoglobin of the annelid Glycera dibranchiata has attracted the attention of several researchers, as it constitutes a readily available source of a monomeric hemoglobin [1,2], in which the usually occurring distal histidine is replaced by an isoleucine residue [3]. Aspects of the molecular conformation have been studied by X-ray crystallography [4,5] and circular dichroism spectroscopy [6]; ligand-binding [7-12] and hydrodynamic properties [1,7] have also been examined. The redox chemistry of this protein has not previously been investigated, and we considered this aspect of its chemistry to be of interest, in light of recent applications of redox methods to the elucidation of mechanistic details in heme protein chemistry [13]. Furthermore, we wished to test Abbreviations: SHE, standard hydrogen electrode; CD, circular dichroism; SCE, saturated calomel electrode; Mb, myoglobin; Hbn, deoxyhemoglobin; HbO2, oxyhemoglobin; CDTA, trans-l,2-diaminocyciohexane-N,N,N',N'-tetraacetate; Hbm, methemoglobin (with ligand indicated); Hbm, monomeric Glycera hemoglobin; HbM, mutant human hemoglobin; Hb^, tetramerichuman hemoglobin.

the applicability to protein chemistry of certain redox approaches used previously [14] in small molecule redox chemistry.

Materials and Methods Live Giycera dibranchiata were obtained from Woods Hole Marine Biological Laboratory * Erythrocytes were washed with 0.15 M Na2SO4/10 mM EDTA and lysed with distilled water. The heavy (oligomeric) and light (monomeric) Hb components were separated by chromatography on Sephadex G-75. The buffer used throughout was (pH 7.0) sodium/potassium phosphate (0.1 M, /t = 0.2) containing 2 mM CDTA (Fluka AG), to sequester transition metal ions which bind to and catalyze the autoxidation of heme proteins [15]. Myoglobin was Sigma type-3 equine, while other reagents were used as supplied by Sigma and Fisher. Hemoprotein was assayed spectrophotometrically as the iron(II) carbonyl, using a molar * Glycera were supplied by MBL as G. americana, but identified as G. dibranchiata by Dr. M. Pettibone of the Smithso-

nian Institute.

0167-4838/85/$03.30 © 1985 ElsevierSciencePublishers B.V. (BiomedicalDivision)

363 absorptivity of 15400 at 540 nm, determined by analysis of total heme using the pyridine hemochrome method (¢ = 32000 at 557.5 nm) [16]. The redox potential measurements were performed under an argon atmosphere (Burdett Oxygen Co., 99.998%) using a 4 cm optical path all-glass flow cell which incorporates an electrochemical cell, in which provision was made for anaerobic reagent addition (calibrated microsyringe) and for the connection of a PAR-173 potentiostat/galvanostat. The three electrodes were: an SCE reference, a Pt-mesh working electrode and Pt-mesh auxiliary electrode, the last isolated in a porous Vycor test tube. Solution potentials were determined vs. the SCE using a Beckman Model 3500 p H / m V meter to monitor a fourth, separate Pt-wire spiral electrode, to avoid problems associated with surface polarization of the Pt-mesh working electrode. Fe(CDTA) 2- solution was prepared by quantitative a d d i t i o n of analytical grade Fe(NH4)2(SO4) 2 • 6H20 to excess H4CDTA dissolved in deoxygenated pH 7.0 buffer. Fe(CDTA)solution was prepared by addition of peroxide to the above, followed by boiling to decompose excess peroxide. Electronic and CD spectra were obtained using Perkin-Elmer 320 and Jasco J41C instruments, respectively. Cyclic voltammetry was performed as described previously [14] using Pt and Hg working electrodes referred to the SCE, the aqueous solution containing 0.1 M KNO 3 in addition to the buffer. Protein redox chemistry was performed at 27°C. Error limits are given as _+ two S.D.s. Direct titrations of Hb m with a variety of reagents [17] were unsuccessful because of instability of potential, associated with the low exchange current of the Hbm/Hbm iii ii couple. Preliminary experiments suggested that the F e ( C D T A ) - / Fe(CDTA) 2- couple [18] was suitable for use as a redox mediator/buffer. The E °' of this couple was checked by cyclic voltammetry to be +0.089 V under our experimental conditions (cf. + 0.090 V [18]. The redox state of the Hbm was monitored in the a/fl-band region, because near-ultraviolet absorption associated with Fe(CDTA)- rendered the Soret region unsuitable for this purpose. The rate of electrooxidation was limited by the maximum anode potential (+0.21 V vs. SCE)

compatible with the Hb~ solution. Above this potential, the formation of HbmO2 evidenced the onset of anodic oxygen generatiofl at the Pt electrode. In practice, anodic preelectrolysis of each Hb~(OH2)/Fe(CDTA )- mixture was found desirable, and, like the anodic oxidation, was carried out at a potential actually not exceeding + 0.185 V vs. SCE. Electroreduction can be effected more rapidly, as the H 2 generated at cathodic overpotentials is innocuous. Results

Addition of a large excess (1.0 mM) of Fe(CDTA) 2- to a solution of aquomethemoglobin, Hbtm"(OH2), effected complete conversion of the protein (6.0.10 -5 M) to the deoxy form, Hb~, indicating that the E °' of the Hbm was more positive than approx. +0.1 V. Addition of Fe(CN)36- as oxidant, in the normal fashion [17,19], resulted in conversion of the Fe(CDTA) 2- and Hb~ to the iron(III) forms. However, the concentration Nernst plot, of log([nbm]/[Hb~]) vs. potential was quite non-linear (Fig. 1), in fact exhibiting a shape which indicated an apparent increase in E °' of the protein as the fraction of oxidized protein increased. E

(mY)

-60

-80

o

-100

•

-120

°°°° o

o

o

oo° -140

_,.o / : o -2.o

-1.5

-1.o

-o.5

o

+0.5

l o g / r Hbmm'~\

Fig. 1. Data from (a) (open circles) Fe(CN)36 --effected oxidation and (b) (closed circles) F e / C D T A - m e d i a t e d redox equilibrium of Hb III m ( O H 2 ) / H b II m at various potentials.

364

Electrochemical oxidation (and reduction) was therefore substituted for chemical oxidation. The Fe(CDTA) -/2- potentiometry was performed initially in the cathodic direction. Fig. 2 shows the effect of cathodic reduction on a solution initially containing 60 #M Hb~i(OH2) and 1 mM Fe(CDTA)-. The observed conversion of Hb~i(OH2) to Hb~ was subsequently reversed by anodic oxidation of the solution. The spectra of Fig. 2 gave concentration data at 567 nm which yielded the linear concentration Nernst plot of Fig. 1. Under redox cycling of the system, the cathodic and anodic data converge, so that a n El~ 2 value of -0.090 +0.008 V vs. SCE (E °' = +0.153 V vs. SHE [20].) was obtained, the plot having a slope (55 mV) close to that expected (59 mY) for Nernstian n = 1 behaviour.

Effect of hexacyanoferrate Observation of Nernstian redox behaviour in the absence of hexacyanoferrate suggested that the anomalous result obtained with Fe(CN) 3- as oxidant was due to this anion or its reduction product. That the latter is indeed the active agent, is

demonstrated by the spectra displayed in Fig. 3. Addition of potassium hexacyanoferrate(II) in increasing concentrations to a solution of Hb~](OH2) caused the solution to become pink, and the absorption spectrum in the a//]-region was characterized by the appearance of two bands of similar intensity, near 575 and 540 nm, which replaced the broader absorption about 500 nm, characteristic of Hb~I(OH2). Circular dichroism spectroscopy of the solution 0.5 M in Fe(CN)64- confirmed that the protein was indeed in HbmO2 form. Similar concentrations of Ru(CN)~- had no effect on the absorption spectrum of Hb~I(OHE). Little oxygenation was apparent when the argon was scrubbed with hexaaquachromium(II) solution, and none in the additional presence of glucose oxidase/glucose (Addison, A.W., Palaniandavar, M., Staphanos, J.J., unpublished data). Generation of HbmOz must be associated with residual dioxygen in the argon stream (0.001-0.004 T0rr).

Potentiometric study of azide binding To further test the applicability of the media-

•

•

/I

u t

2.0

\

o

I-"

t

~,

2.c

j

1.E 1.0

x

Mot

:

I '~ ,' I

Hb/'~ /

1.0

0.i / 0.5

400

500

600

700

k(nm) Fig. 2. Absorption spectra obtainedduring Fe/CDTA-mediated cathodic reduction of Hbm(OH2) (spectrum-I) to Hbu (spectrum-8).

I

I

400

500

6(~0 ~(nrn)

, 70O I

Fig. 3. Spectra resulting from treatment of Hb~](OH2) with Fe(CN)64-: ( . . . . . . ) Hb~I(OH2) with no added Fe(CN)t4-; ( ) with 0.05 M Fe(CN)t4-; ( ) with 0.5 M Fe(CN)64- and ( - - - - - - ) spectrum of Hbm(O2) for comparison. The higher Fe(CN)~- concentration corresponds to the higher, longer-wavelength Soret absorption.

365 tor/buffer method to the determination of heme ligation equilibria, we investigated potentiometrically the interaction of the H b ~ / H b ~ couple with a model ligand, N 3. Azide is known to bind strongly to iron(Ill) heroes, but not significantly to irOn(II) hemes, and the thermodynamics of Glycera Hb~l(N3) formation has been studied previously, by Seamonds et al. [7] and Beetlestone et al. [11]. Increasing concentrations of NaN 3 were introduced into a solution containing the Hb m m/Hbmn and Fe(CDTA)-/Fe(CDTA) 2- couples at equilibrium. The initially totally iron(III) solution had been adjusted to a potential of -0.146 V. vs. the SCE by cathodic reduction of Fe(CDTA)-, and the solution containing mostly reduced protein was thus redox (potential) buffered by the Fe(CDTA)-/2- couple present in large excess over the Hbm. Cyclic voltammograms of Fe(CDTA)are unaffected by addition of 0.4 M azide. Addition of azide effects net conversion of Hb~ to Hb m, the latter manifested spectroscopically mostly as Hblm"(N3). The plot of ([Hbmln]/[Hbm])l! vs. [Nf] in Fig. 4 was obtained using absorbance data near 640 rim, where there is an isosbestic point for the Hb~i(OH2) and Hblmll(N3) spectra. The resulting value of log K, for substitution by N 3 of the water molecule in Hb~l(OH2) is 2.98 +_

6.0

[Hb~] [Hb~]

4.0

/.

2.0 0.0

• I 1.0

I 2.0

I

3.0

I 4.0

I 5.0

6.0

[N;] X I0-a Fig. 4. Data for spectropotentiostaticdeterminationof azide bindingat ionicstrength0.2 M.

0.08. Beetlestone et al. reported log K = 4.29 at pH 6.2, /~ = 0.5 [11], while the work of Seamonds and co-workers [7] indicates rather lower values, of 2.57 (pH 7.3) and 2.66 (pH 6.45, both in 0.1 M phosphate). To check our potentiostatically obtained result, we also measured Hb~l(N3) formation by the usual spectrophotometric procedure, involving addition of N f to Hb~I(OH2) in 0.1 M buffer. The plots obtained, of [Hb~I(N3)]/ [Hb~I(OH2)] vs. [N3] were not linear over the full complexation range, and the slope of a typical log plot was equivalent to a Hill coefficient of n = 0.643, with log K = 2,25. Howev6r, the required limiting spectrum (more than 99% Hb m N 3) is thus associated with [Nf] > 0.4 M, a pbssibly significant change in ionic strength during the titration. We therefore titrated Hb~l(OH2) with NaN 3 in 1.0 M phosphate buffer, and the resulting data, plotted as log(a/1 - a)'* vs. log[N3] gave n = 0.94 and log K = 3.80 +_0.52 via interpolation to [Hb~(N3)]/Hb~(OH2)I = 1.00. Discussion

Glycera Hb E °' The redox potential determination using the Fe(CDTA) -/2- couple relies on~ an approach slightly different to that used previously [18], as we elected to measure the solution potential directly, rather than deducing it from the concentrations of the added Fe(CDTA)- and Fe(CDTA) 2-. In addition, the method yields the advantages, that the total CDTA species concentratiofi remains constant, and there is no need to correct absorbances for the dilution effects of reagent solution addition. The E °' value obtained is s i ~ l a r to that of several other hemoglobins [21-24]~'under similar conditions, but thus quite different :from those for myoglobins and erythrocruorins [2,1,25] with one important exception, namely Aplysia Mb, which also lacks the distal histidine [21]. The active site local environment is thus more important in controlling the iron chemistry than are" the perturbations associated with quaternary structure, though these are significant [13]. In c o m p ~ s o n with the other monomeric systems, such as Mb [21] and * a, fractionin complexedform.

366 Hb~ [23], substitution of distal histidine by a nonpolar residue destabilizes the aquo-met state by approx. 100 mV (10 k J), which is of the right magnitude to reflect an H-bonding interaction between histidine and coordinated H 2 0 in the 'normal' proteins. The results may be placed in a wider context by Fig. 5. It has been proposed [19] that oxidation and oxygenation thermodynamics may be linked via RT E °'= E* + --~ In Pl/2 (wherein E* refers to an arbitrary heme with Pl/2 = 1 Torr), and indeed, oxidation closely parallels oxygenation over a free-energy range of about 8 kJ in the chemistry of Hb A [13,36]. The slope (66 + 16 mV) of the regression through points 1-12 ( E °' = 96 + 66 log P~/2) is not significantly different from the value (59 mV) in the above relationship, so that Fig. 5, covering a range of about 22 k J, is the strongest evidence to date, for the generality of the oxygenation/oxidation parallel. It also signals the uniquely structured active site cavity in leghemoglobins. The increased oxygen affinity associated with increased reducing tendE o'

(mV)

e14

+200

e13

e6

~ 5Z l o9l l e8

+100

el

e4

o10

e3

--100

I --2

~ --1

I 0

I +1

t +2 Log p ,/.

Fig. 5. Redox potentials of hemoglobins and myoglobins mapped against oxygenaffinity.The data are for : (1) Aplysia limacina Mh [21,26]; (2) Physeter macrocephalus Mb [21,26]; (3) Candida mycoderma Hb [27]; (4) Chironomus thummi Hb [21,28l; (5) HbM-HydePark [22,29l; (6) HbM-lwate[22,30]; (7) Hb,, [23,31,32]; (8) Hbp [23,31,32]; (9) HbF [23,24]; (10) Hbr [23]; (11) HbA [22,24,33];(12) G. dibranchiata H b m (this work and Ref. 9); [13] Leghemoglohin-a,and (14) Leghemoglobin-c [34,35]. The Pi/2 values are referred to a standard state of 1 Torr.

ency of the iron(II) state parallels the correlation in behavior observed between oxidation and oxygenation of cobalt(II) complexes [37]. These observations have been interpreted in terms of electron-transfer toward 02 during Fe-O 2 bond formation [38,39]; the alternative approach [40] focuses on the spin-pairing between Fe and 02. The two viewpoints are perhaps united by consideration of the role of the iron(II) porphyrin H O M O energy. There is also some correlation (r--- +0.57) between the 02 and CO binding constants [41] of heine proteins. However, this is observably weaker, supporting the idea [42] of reduced importance of net ground state d(Fe)-~r*(CO) charge transfer, despite the apparently still polar nature of the F e - C O linkage [43,44]. In addition, this latter correlation is markedly affected by distal influences; CO binding is considerably enhanced vs. 02 binding for Hb~ [10], in which the usual destabilization of the carbonyl by distal histidine is absent. Hexacyanoferrate binding The non-Nernstian redox results, as well as the spectra in Fig. 3, are interpretable only in terms of quite strong binding of Fe(CN) 4- to Glycera hemoglobin. The positive cooperativity between Fe(CN) 4- binding and oxidation is paralleled by enhancement of oxygen affinity in the presence of Fe(CN) 4-. Fig. 3 shows (and the CD confirms) that the Fe(CN)64- does not coordinate to the heme iron, but leads to generation of HbmO2 from Hbl~ I under conditions where no HbmO2 is observable in the absence of Fe(CN) 4-. That is, the oxygen affinity of Hb~ is enhanced by Fe(CN)64--binding *, and the consequent removal of Hb~ from solution as HbmO2 perturbs the effective E °' of the Hb m / H b m ii couple, so that Hb~I(OH2) is reduced by Fe(CN)64- . These consequences are perhaps surprising, as Fe(CN)64- is far too weak a reductant ( E ° ' = +0.43 V) to be expected to reduce Hb~I(OH2) ( E ° ' = +0.15 V) extensively in the normal fashion **. These results * Direct enhancementof 0 2 affinity by hexacyanoferrate(II) addition is observedas well. ** The potential is not defined unless hexacyanoferrate(III)is also present; after Hbm reduction, E would be about + 0.22 V.

367 moreover indicate that the use of Fe(CN)36 - as a redox titrant is better avoided. Binding of Fe(CN)64- to heme proteins has been observed previously, for tetrameric and polymeric hemoglobins-[13,25]. Because of the wellcharacterized allosterically effective binding of polyanions to the subunit interfaces of tetrameric H b [45], one might have anticipated that the monomeric Glycera Hb would lack a strong binding site for anions. This is clearly not the case, nor does a stereochemical allosteric mechanism necessarily apply for Glycera hemoglobin. In addition, the binding of polyanions to H b A raises the Ell 2 [36,46] and decreases the oxygen affinity [36,47]. Glycera H b m is thus noteworthy, inasmuch as the anion effect of Fe(CN) 4- is the opposite of that usually observed for hemoglobins.

Azide binding Comparison of our azide-binding data with those available from previous workers [7,11] leads us to suspect that the principal sources of apparent variation in the reported log K values lie (a) in the method of data treatment and (b) the differences in ionic strength. The value of log K obtained by the spectropotentiostatic procedure is likely representative of the situation at the lower ionic strength. The approach relies on methodology we have used previously [14] in small-molecule redox chemistry, where the usual approach is to determine the effect on the system mid-point potential, E~/2, of the added ligand, L, which binds to the protein metal: Hbm+ xL ~ HbmL:, For heme proteins, it is rare that x > 1, so that for x=l: [HblmII] [Hbn ]

eQl+gm[L] 1+ K'I [L]

where H b u = total iron(II); ---m l-lhm = total iron(Ill); K iI and K m are the 1 : 1 ligand binding constants for these two oxidation states, and:

e ffi (e.~o,o,~oo- e ° ' ) n F / R r E °' referring to 'unligated' protein. For azide,

K I! < 1 << Kill( = K), so: [ Hbu] [Hb~']

eo+eOK[L]

In the presence of the redox-buffer, e Q is constant, so that the response of [Hb m ]/[Hbm] i1 to added ligand can be determined conveniently by the absorption changes at a wavelength which is isosbestic for HblmI[ and H b m L. A further advantage [14] of the method is that the limiting spectrum of fully formed Hb m ill L is not required, which makes it particularly gainful in the determination of low K values, absorbance corrections, for the volume of ligand solution added, can thus be kept smaller than those in the direct spectrophotometric method. In the conventional spectrophotometric approaches, the value of either n or elim must be guessed if ~lim is not available, and correction of [L] for bound ligand is not always possible [48-51]. Graphical data treatments reveal the integral value(s) of n in the spectropotentiometric method, and allow correction for bound ligand, References

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