formationand stabilitiesof free bilirubinand - clinical chemistry

III

CLINICAL CHEMISTRY, Vol. 17, No. 12, 1971 1165

Formationand Stabilitiesof Free Bilirubinand BilirubinComplexeswith Transitionand Rare-EarthElements

Rance A. Velapoldi and Oscar Menis

We investigated complexes formed between bilirubin and transition orrare-earth elements, and their relative stabilities.Relativeratesof complexformation were analogous to metalloporphyrin and metal complex stabili-ties, according to the Irving-Williams series. Transition metals that formedstrong, square planar complexes caused rapid bilirubin degradation. Iron-(II) was oxidized to iron(IlI) in the bilirubin complex. Of the Ianthanidecomplexes, only samarium(llI) showed covalent bonding tendencies.Comparative molar absorptivities were calculated for several of the metal-bilirubin complexes. The stability of bilirubin in several solvents and undervarious experimental conditions is reported. At room temperature, biIirubin photooxidizesto biliverdin on exposure to laboratory light.

Additional Keyphrases bilirubin degradation #{149} spectral shifts #{149} solventeffects #{149} lanthanides photooxidation

BR’ is unstable in solution; 1% degrades perday in chloroform, 40% in base. In the body,bilirubin is produced by the metabolic breakdownof hemoglobin (1-3). It is a linear tetrapyrroleconsisting of two conjugated dipyrrole segmentsjoined by a labile methylene bridge. It can existin several tautomeric forms, two of which are thebis-lactam, I, and bis-lactim, II, structures (4-8).

Me R Me P P Me Me R

OoH H H H

R- -CH=CH,

P - -CH2CH2COH

Me R Me P P Me Me R

IIR- -CH=CH2

P - -CHCH2CO,H

Because of its chemical characteristics, it wasthought that BR existed mainly in the bis-lactamform (1, 4-8). However, in 1962, O’Carra (9) and,

From the Institute for Materials Research, National Bureauof Standards, Washington, I). C. 20234.

Nonstandard abbreviations used; ne, bilirubin; uvo, bihver-

din; EDTA, ethylenediaminetetraacetjc acid EtOH, ethanol;MeOH, methanol.

Received May 26, 1971; accepted Sept. 2, 1971.

in 1965, van Roy and Heirwegh (10) reported

briefly that BR and meso-BR form relatively stablecomplexes when zinc ions in alcoholic solution areadded. Fog and Bugge-Asperheim (11) haveshown that adding zinc ions to an alkaline solutionof BR results in BR degradation, while addition ofEDTA stabilizes the BR solution.

Considering the proposed hydrogen-bound struc-ture for urobilins (8, 12) and the observed BR proto-tropy (8, 13, 14), it is probable that BR exists insolution as a hydrogen-bound mono-lactam-mono-lactim structure, (III).

The circular configuration and nucleophilic nitro-gen now present are both conducive to complexformation, analogous to the metalloporphyrins(15), which in the absence of complex ions mayexist partially in solution in the hydrogen-boundform (III).

Metal’ Peakb, nm

1166 CLINICAL CHEMISTRY, Vol. 17, No. 12, 1971

To elucidate BR systems more fully, we under-took a study to determine the effects on BR sta-bility of adding metal ions and of various otherexperimental conditions.

Materials and Methods

Equipment2

A spectrophotometer (Cary Model 14 Record-ing Spectrophotometer, Cary Instruments, 2724 S.Peck Rd., Monrovia, Calif. 91016) equipped withconstant-temperature (± 0.1#{176}C)cell block wasused to record the ultraviolet and visible spectra.A curve resolver (Dupont 310, E. I. du Pont deNemours, Instruments Products Division, Wil-mington, Del. 19898) was used to differentiate thevarious species in solution if absorbance curvesoverlapped. Use of a nitrogen manifold andSchlenk (16, 17) equipment ensured that air wasexcluded.

Reagents

All reagents and solvents were reagent gradeand not further purified except for those solventsused in the inert atmosphere tests, which were dis-tilled under purified nitrogen that had been passedthrough a sodium-potassium alloy. The lantha-nide acetates were at least 99.9% pure.

Methods

Solvent systems. The following solvents were in-vestigated to determine the system most conduciveto complex formation: chloroform, absolute etha-nol :chloroform (1:1, by vol), methanol: chloroform(1:1, by vol), tetrahydrofuran, dimethylsulfox-ide, and potassium cyanide-dimethylformamide(18). For example, 0.05 ml of samarium(III)acetate (2.42 mmol/liter) was added to andmixed with 1.3 ml of a BR solution (4.6 imol/liter)in ethanol: chloroform, resulting in a 4:1 molarratio of Sm(III) to BR.

We recorded absorption spectra at 25.0#{176}Cas afunction of time to determine the spectral shiftsattributable to complex formation. In some ex-periments, after spectral shifts were completed,strong coordinating higands such as EDTA or 1,2-diaminocyclohexanetetraacetic acid were added.The Fe(III) was reduced by using a reducing agentsuch as sodium dithionite or hydrazine.

Stabilities. The more stable BR and metal-BRcomplexes were placed in Schlenk-equipped

In order to describe experimental procedures adequately.

it is occasionally necessary to identify commercial productsand equipment by the manufacturer’s name or label. In noinstance does such identification imply recommendation orendorsement by the National Bureau of Standards, nor does itimply that the particular product or equipment is necessarilythe best available for that purpose.

matched quartz and Pyrex spectrophotometriccells and spectra were recorded intermittentlyover long periods of time to determine solution sta-bilities. Solutions in these studies remained inthe dark under refrigeration or under other experi-mental conditions listed in Table 3.

Results and Discussion

Complex Formation

Transition metals. Added metal ions, except so-dium, produce one of two results: a bathochromic(toward the red) or hypsochromic (toward theviolet) shift (19) of the BR peak at 454 nm,followed by the disappearance of the new peak,or there is a disappearance of the BR peak withoutdefinite new peak formation. These shifts to-ward the red and violet suggest complex forma-tion and not simply salt formation with the propio-nic acid functional group, because salt formationwould cause shifts in essentially one direction.Spectral shifts are best illustrated by the additionof samarium(III) to BR, Figure 1. Complex for-mation in this case does not appear to be compli-cated by side reactions such as oxidative degra-dations as observed with transition metals becausethe Sm(III)-BR complex spectrum does not dis-appear rapidly. This particular shift requires 12h for completion. The observed isosbestic pointat 414 nm suggests the presence of only two spe-cies in solution-the free and the complexed BR

(20). Table 1 summarizes the results of metal

Table 1. Spectral Shifts upon Metal Ion Additionto BR Solution

Molarabsorptlvlty,

(lIter mofr’Stable’ cm ‘)

BR 454 yes 61,100Zn(ll) 476 noCu(ll) (1)420,(2)352(1)no,(2)noNi(lI) 476 noCo(ll) 423, 470(B)4 noFe(ll) (1) 424, (2) 463 (1) no, (2) yes 53,300[for

Fe(l1)1Fe(lll) 450(B)4 no 51,300Er(lll) 385 noGd(lll) 385 noPr(lll) 385 noEu(lll) 387 noNd(ll) 391 partially 56,600Sm(lll) 395 yes 46,400Ca(ll)Na(l)

464, 495 (S)6no shift

[M’i/[BR] = 4.b Maximum absorbance of (1), initial peak and (2), final peak.‘Stable means less than 3% absorbance decrease during

40 days.d (B) = broad.

(S) = shoulder.

ILlC..)2‘3CDa:0‘/)CD‘3

I . BILIRUBIN2. oSm(ffl) AFTER 5mm3. AFTER 45mm4. 90mm5. 1005 mit, (EOUIL.VALUE)

.0 I I

FeUD-BR -0.9 COMPLEX ,_#{149}-‘ Fe(ilI)6R

/ ......- CLEX -

U .. \:‘.‘o\’. -

0.6 #{149}‘./),‘/ / \..“ \#{149}‘

0.5 /‘7/ //CD -fi /‘304#{149} 4i .‘

1/’ / --- SR ALONE \ C:9 -0.3 ;Y.i/’ .: 25n,n \/ 05+ F,lfl>to+ 55n

-. - - 000Fe(Ulto+ 8O,nn \02 --- BR+FeIU}tot- I2O,on \..

-- ----- 05* F, (III to +4,n,n \\. \\

\.. s”.0.1- I I I

500


400 450X, NANOMETERS

Fig. 1. Complex formation: Addition of samarium-(Ill) acetate in ethanol to bilirubin in chloroform:ethanolresults in a slow blue spectral shift

ion addition and indicates the positions of absorp-tion peaks and their stabilities towards degra-dation. The positions of the absorption bands forthe complexes between BR and transition metals(except copper) follow closely those of the metallo-porphyrins (15). As electrons are added to themetal “d” orbitals [as in Zn(II), Ni(II) and Co-(II)], electron repulsion creates a greater electrondensity at the periphery of the ring, and shifts tolonger wavelengths are observed.

Qualitatively, the rate of decrease in BR absorp-tion caused by metals is in the order Cu(II) >Ni(II) > Co(II) > Zn(II), which is the order forthe relative rates of formation of the metallopor-phyrins and for the stability of metal complexesaccording to the Irving-’Williams series (21).An exception to the above order is observed uponFe(II) addition, which should come after Zn(II)for rate of BR absorption decrease but comes be-fore Cu(II). The immediate shift (within 2 mm)of the 454 nm peak to 424 nm is followed by aslower (4 h) red shift. After 4 h, addition of thereducing agent, sodium dithionite, results in ashift reversal-the 463 nm peak decreases and the424 nm peak increases, showing that the firstpeak formed at 424 nm is attributable to an Fe-(11)-BR complex and the subsequent shift is at-tributable to Fe(II) oxidation to Fe(III) whilecomplexed to the BR. Similar redox reactions areobserved in the Fe-porphyrin systems (22-24).A proposed sequence for these reactions is as fol-lows:

(0)BR + Fe(II) Fe(II)-BR complex ..____

(454 nm) ETA, H * (424 nm) dithionite

Fe(III)-BR complex (1)(463 nrn)

The isosbestic point in the spectra during the Fe-(11)-BR complex oxidation supports the presenceof only two species in solution, Fe(II)- and Fe-(111)-BR complexes (Figure 2). Overlap of thetwo absorption peaks makes it difficult to deter-

400 450 500WAVELENGTH, nm

Fig. 2. Complex formation and iron(ll)-BRoxidation toiron(lll)-BR: Iron(ll) addition to a bilirubin solutionforms an iron(ll)-BR complex followed by iron(ll) oxida-tion to iron(lll)

mine rate data; however, resolution of the bandsby use of a curve resolver gives results which, whenplotted, adhere to the expected plot for a second-order reaction in the oxidation step.

Cu(II) addition results in rapid complex forma-tion with subsequent degradation. This may beexplained on the basis of formation of a squareplanar or distorted square planar configurationthat adds ring strain at the methylene carbon,causing increased lability and thus oxidation atthe periphery of the ring. Similar oxidations arefound in metalloporphyrin systems (25). In addi-tion, copper is also a very good oxidizing metal inbiological systems (26, 27). The other metals formtetrahedral complexes to some extent, and so strainat the methylene carbon is decreased.

Lanthanide metals. Various lanthanides addedto BR solutions produced similar hypsochromicshifts. The small red shifts in the complexesbetween BR and praseodymium, neodymium, orsamarium (385 to 395 nm) have been reportedpreviously for the anhydrous lanthanide salts inother systems (28, 29). The shift was explainedby Ephraim and Bloch (29) as an expansion ofthe lanthanide orbital radius owing to increasedcovalency. This explanation was designated anephelauxetic effect by J#{248}rgensen(30). Tsutsuiand Gysling (31) recently invoked similar reason-ing for increased covalency in lanthanide com-plexes. In their work, only Sm(III) showed co-valent bonding tendencies among several newlanthanide indenides, which is consistent with ourfindings. This small shift may be interpretedsimply as a withdrawal of electrons from the ligandring system by the lanthanide ion.

Spectral shifts caused by the addition of Pr(III),Er(III), Eu(III), and Gd(III) were rapid (equi-librium values reached within 20 mm). On addi-tion of Sm(III) and Nd(III), equilibrium was

Table 2. The Effect of M’ Concentration on Com-plex Formation

Metal LM’I/(BR] Peak 1, nm

Zn(ll) 100 4924 476

Fe(ll) 4 4241 445 (B)#{176}

Co(ll) 10 420, 475 (B)’4 422, 471 (B)1 425, 468 (B)

Cu(ll) 100 4184 4211 4250.1 no shift

Sm(lll) 10 3954 3951 401

‘(B) = broad.Ethanol and chloroform in equal volumes was the solvent in

all cases.


slowly attained (15 and 3 h, respectively); La-(III) addition caused a slight shift with bandbroadening, but no evidence of actual complexformation.

After equilibrium times were reached in the for-mation of the Sm(III), Fe(III), and Nd(III) com-plexes, either 1 ,2-diaminocyclohexanetetraaceticacid or EDTA was added, resulting in a reversal ofthe observed spectral shifts. After about 90 mm,the free BR spectrum, with slight broadening butsimilar absorbance readings, was observed. Thisproves metal ion-BR complexation and not BR oxi-dation to BVD followed by metal complexation withthe BVD, a reaction that is known to occur (32).Comparative molar absorptivities for the Sm-(III)-, Nd(III)-, Fe(II)-, and Fe(III)-BR com-plexes calculated relative to free BR are listed inTable 1.

Metal Ion Concentration Effect

Metal salt solutions were prepared and added tothe BR solution resulting in different [M” + / [BR]

ratios. Table 2 summarizes the results of themetal ion concentration effects. A ratio of 4 re-sulted in essentially 100% complex formation.A 4:1 molar ratio of Zn(II) : BR gave a batho-chromic shift to 476 nm, in agreement with earlierreports (9, 10); however, when a large excess ofZn(II) ions was added, there was a further redshift to 492 nm. Increased interactions by excessZn(II) ions with the pi electron density in therings may withdraw electron density toward theperiphery and cause the further bathochromicshift. This phenomenon was not observed withthe other metals.

Stability Study

Solvent systems. The stability of BR under vari-ous experimental conditions is summarizedTable 3. It appears that a mixture of equal vol-umes of ethanol and chloroform stabilizes BR lfl

solution: its concentration decreased by only0.051% per day. BR 1S less soluble in this solventmixture than in chloroform alone. It is interest-ing to note that a sample of very nearly pure BR,3

dissolved in chloroform and kept from light, wasalso highly stable [Deardorif, E. R., private com-munication I (no measurable decomposition in oneweek). This observation, along with Fog’s pre-vious paper (11), suggests that metal impurities(especially transition metals) increase the rate ofBR degradation. Complexation of the metal im-purities with EDTA resulted in stabilization of theBR solutions. From our work, it is also evidentthat complexation of the BR with metals that form“covalent” complexes [Sm(III) or Fe(III)] alsoresults in stabilization.

As can be seen, equal volumes of methanol andchloroform is a poor solvent for BR. Ascorbic acidaddition did not inhibit BR oxidation as expected,but rather increased it.

Metal complexes. Only Fe(III) and Sm(III)gave complexes with BR that were considered to bestable. Zn(II), in contrast to earlier reports(9, 10), did not give a stable complex, although“stability” in this paper means degradation of lessthan 3% (measured spectrophotometrically) dur-ing a 40- to 60-day period.

The slower rate of decomposition (spectral banddisappearance) for the Sm(III) complex does notappear to be significant when compared to the de-crease for BR itself. The Nd(III) complex de-grades 10-fold faster than the Sm(III)-BR complex.The other lanthanide complexes decreased quiterapidly (less than 1 or 2 h for spectrum disappear-ance). Increased covalency in the following order,Sm(III) > Nd(III) > lanthanides may be the rea-son for the decreased degradation rates. In addi-tion, flocculent material was noted in the lantha-nide complex solutions, except for Sm(III) andNd(III). This flocculent material has not yetbeen identified; however, if the solution is shakenthe metal-BR complex spectrum reappears, whichsuggests that the flocculate is the precipitated com-plex. The resulting absorbance, upon shaking, is10-fold less than was the original absorbance forthe complex.

Effect of light and temperature. A limited studywas made to determine the effect of temperatureand light on the rate of BR decomposition. Solu-tions of BR in EtOH : CHC13 (solutions 4, 9, 10,

‘Sample obtained from Dr. Robert Shaffer, the OrganicChemistry Section, Analytical Chemistry Division, NationalBureau of Standards, Washington, D.C. 20234.

Table 3. Free and Complexed BR Stabilities UnderVarious Experimental Conditions

Lengthof

test,days

C-,2oxm600)ox

Fig. 3. Photodecomposition of bilirubin to biliverdinoidpigment, followed spectrophotometrically as a functionof time

U’

CD

uJ

I-

Fig. 4. Determination of bilirubin photodecompositionto biliverdinoid pigment by resolving individual curves ofFigure 3. Corrected areas under respective curves areproportional to relative concentrations

(2) We are pleased to acknowledge the partial support of thiswork by the National Institutes for General Medical Sciences,NIH.


No. Solvent Add ItIvea

1 MeOH:CHCI,

2 MeOH:CHCI, ascorbic acid3 EtOH:CHCI, in air4 EtOH:CHCI, under N,5 EtOH:CHCI, ascorbic acid6 EtOH:CHCI, + Sm(lll)7 EtOH:CHCI, + Nd(llI)8 EtOH:CHCI,+ Fe(Il)9 SoIn 4, refrigerated

10 Soln 4, at room temp, in dark11 SoIn 4, at room temp, in light

‘As measured by absorbance decrease.

DegradatIon p.r

day,’ %

5 particulatesseen

40 0.62540 0.05163 0.05340 0.08663 0.04556 0.49560 0.05146 0.01310 0.3014 24.738

and 11, Table 3) were used in this study. Duringthe test period, solution 4 was kept in the darkunder refrigeration between measurements. Ab-sorbance measurements were made often, necessi-tating removal from the refrigerator followed bywarming to 25#{176}C.During testing, solution 9was kept refrigerated and was not disturbed dur-ing the 46-day test period. This solution showed4-fold less decomposition than solution 4. Solu-tion 10, brought to room temperature but not ex-posed to light, degraded more rapidly than solution4, which was exposed to room temperature onlyintermittently. Solution 11, on the other hand,showed total decomposition in only 4 days uponexposure to light at room temperature.

Decrease in the BR absorption band during this4-day period was accompanied by the appearanceof new peaks at 377 and 700 nm, with a shoulderat 655 nm (Figure 3). These peaks are attribut-able to biliverdin formation (identified by the finalabsorption spectrum), which must form by aphotochemical reaction. Separation of the BR

and BVD absorption bands by use of the curve re-solver resulted in data which, when plotted aspercent area under curve (relative concentration)vs. time, yielded the data presented in Figure 4.BR decreases as fast as BVD appears (equal slopes).Molar absorptivities calculated for the BVD fromthe spectra agree quite well with those given in theliterature (33, 34): experimental, 377 nm, e = 42,-200 liter mo11 cm’, 665 nm, e = 15,540 litermo!’ cm’ as compared to the literature values of41,700 and 13,400 liter mol’ cm, respectively.

The zero-order straight-line plot for these datasuggests a photochemical reaction sequence such as

(a) A+hv-+A’

(b) A*+BR#{247}Bvn+A

where Equation (a) is the limiting step, and A issome species in large excess.

In summary, we found that BR forms complexeswith various metals, resulting in either a varyingdegree of stability or rapid BR disappearance (deg-radative oxidation). Metals that form strong,square planar complexes oxidize BR more rapidlythan other metals, probably by inducing ringstrain at the already labile methylene bridge.Addition of EDTA to stabilize BR solutions (11) un-doubtedly removes the degradative effects of tran-sition metals by complexation. Light causesphotooxidation of BR to BVD.


References

1. Lemberg, It., and Legge, J. W., In Haematir& Compounds and

Bile Pigments, Interscience, New York, N.Y., 1949.

2. With, T. K., In Bile Pigment.s, Academic Press, New YorkN.Y., 1968.

3. Pullman, B., and Perault, A.-M., On the metabolic breakdownof hemoglobin and the electronic structure of the bile pigments.Proc.Nat.Acad. Sci., U.S. 45, 1476 (1939).

4. Fischer, H., Pleininger, H., and Weis.sbarth, D., Bile pig-ments (XXX). Structure of bilirubin and on bilirubinoid pig-ments. Z. Physiol. Chem. 268, 197 (1941).

5. Fischer, H., and Pleininger, H., Synthesis of biliverdin andbilirubin. Z. Physiol.Chem. 274, 231 (1942).

6. Gray, C. H., and Nicholson, D. C., The chemistry of the bilepigments. Structures of stercobilin and d-urobilin. J. Chem. Soc.,London, 3085 (1958).

7. Gray, C. H., Kulczycka, A., and Nicholson, I). C., The chemis-try of the bile pigments. Part III. Prototropy of bilirubirt to averdinoid pigment. J. Chem. Soc., London, 2268 (1961).

8. Gray, C. H., Kulczycka, A., and Nicholson, D. C., The chemis-try of the bile pigments. Part IV. Spectrophotometric titrationof the bile pigments. J. Chem. Soc., London, 2275 (1961).

9. O’Carra, P., Zinc complex salt formation by bilirubin andmesobilirubin. Nature 195, 899 (1962).

10. van Roy, F., and Heirwegh, K., Ability of bilirubin to formzinc complexes. Arch. mt. Physiol. Biochim. 73, 535 (1965).

11. Fog, J., and Bugge-Asperheim, B., Stability of bilirubin.Nature 203, 756 (1964).

12. Cole, W. J., Gray, C. H., and NichoLson, D. C., The chemis-try of the bile pigments. Part V. The stereoisomerism of uro-bilins. J. Chem. Soc.,London, 4085 (1965).

13. Nichol, A. W., and Morell, D. B., Tautomerism and hydro-gen bonding in bilirubin and biliverdin. Biochim. Biophys.Acta177, 599 (1969).

14. Badger, G. M., Harris, R. L. N., Jones, R. A., and Sasse,J. M., Porphyrins. Part I. Intramolecular hydrogen bonding inpyrromethanes and porphyrins. J. Chem. Soc., London, 4329(1962).

15. Falk, J. E., Porphyrins and Metallopor phyrins, Elsevier,New York, N.Y., 1964.

16. Schlenk, W., and Thai, A., Uber Metailketyle, eine grosserKiasse von Verbindungen mit dreivertigen Kohlenstoff. Ber.Dent. Chem. Ges.46, 2840 (1913).

17. Schick, W. H., A new approach to handling air-sensitivecompounds. Amer. Lab., June, 49 (1970).

18. Lemberg, R., and Legge, J. W., Haematin Compounds andBile Pigments. Interscience, New York, N.Y., 1949.

19. Brasted, it. C., and Cooley, W. E., Physical methods in co-ordination chemistry. In Chemistry of the Coordination Com-pounds, Bailar, J. C., Jr., and Busch, D. H., Eds. Reinhold, NewYork, N.Y., 1956, p 565.

20. Stranks, II. It., In Modern Coordination Chemistry, Lewis,J., and Wilkens, it. G., Eds. Interscience, New York, N.Y.,1960, p 84.

21. Irving, H., and Williams, R. J. P., Order of stability of metalcomplexes. Nature 162, 746 (1948).

22. Wang, J. H., Nakahara, A., and Fleischer, E. B., Hemo-globin studies, I. The combination of carbon monoxide withhemoglobin and related model compounds. J. Amer. Chem. Soc.80, 1109 (1958).

23. Tsutsui, M., Velapoldi, R. A., Hoffman, L., Suzuki, K.,and Ferrari, A., Unusual metalloporphyrins (III). Inducedoxidation of Co(II) and Fe(II) porphyrins by unsaturated hydro-carbons. J. Amer. Chem. Soc. 91, 3337 (1969).

24. Whitten, D. G., Baker, E. W., and Corwin, A. H., Spectralstudies of the “d6” metalioporphyrins. Ligand and solvent fieldeffects on ferrous and cobaltic mesoporphyrin. IX. J. Org. Chem.28, 2363 (1963).

25. Castro, C. E., and Davis, H. F., Peripheral attack in thereduction and oxidation of iron porphyrins. J. Amer. Chem. Soc.91, 5405 (1969).

26. De LaMarc, H. E., Kochi, J. K., and Rust, F. F., Oxidationand reduction of free radicals by metal salts. J. Amer. Chem. Soc.85, 1437 (1963).

27. Haines, R. M., and Waters, W. A., Properties and reactionsof alkyl free radicals in solution. Part VIII. The reducing actionof some water soluble radicals. J. Chem. Soc., London, 4256(1955).

28. Sinha, S. P., In Complexes of the Rare Earths, PergamonPress, New York, N.Y., 1966.

29. Ephraim, F. and Bloch, R., Uber die Kontraktion beiVerbindungshildung. Ber. Dent. Chem. Ges.59, 2692 (1926).

30. J#{216}rgensen,C. K., In Progressin InorganicChemistry, 4,

Cotton, F. A. Ed. Interscience PubI., New York, N.Y., 1962.

31. Tsutsui, M. and Gysling, H. J., Evidence for a covalentorganolanthanide. Tris(indenyl)samarium-THF. J. Amer. Chem.Soc. 90, 6880 (1968).

32. OhEocha, C., In Chemistry and Biochemistryof Plant Pig-ments, Goodwin, T. W., Ed. Academic Press, New York, N.Y.,1965.

33. Bonnet, R., and McDonagh, A. F., Oxidative cleavage ofthe haem system: The four isomeric biliverdins of the IX system.Chem. Commun. 237 (1970).

34. Gray, C. H., Kulczycka, A., Nicholson, D. C., and Petryka,Z., The chemistry of the bile pigments. Part II. The prepara-tion and spectral properties of biliverdin. J.Chen. Soc.,London,2264 (1961).

formationand stabilitiesof free bilirubinand - clinical chemistry

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