Bilirubin-Metal Ion Complexes John D. Van Norman and Michelle M. Humans Department of Chemistry. Youngsto wn State University. Youngsto wn. Ohio 4 4 5 0 3
Velapoldi and Menis ( I ) have shown t h a t bilirubin, a bile pigment, forms definite, strong metal-ion complexes in chloroform-ethanol mixtures and have studied the spectral shifts of the bilirubin absorption as a function of added metal ion. They concluded that bilirubin complexes with a variety of metal ions and those metal ions which form strong square planar complexes promote oxidation of bilirubin by placing a strain on the bilirubin structure. Van Norman ( 2 ) has studied the electrochemical oxidation of bilirubin in N,N-dimethylformamide and found that controlled potential coulometry could be used as a possible alternate method of checking the purity of bilirubin. This present investigation is concerned with the electrochemical behavior of bilirubin and bilirubin-metal ion complexes in aqueous media. Since bilirubin is fairly insoluble in acidic media, all measurements were made in a phosphate buffer of pH 7.8 or 8.0. The use of this pH also minimizes the problem of intra- and intermolecular hydrogen bonding interfering with complex formation. EXPERIMENTAL E q u i p m e n t ; All electrochemical measurements were obtained
with a National Instrument Laboratories “Electrolab” ( 3 ) ,a multifunctional electroanalytical system in conjunction with a Valtec Model 1024 X-Y recorder. All p H measurements were taken using either a Chemtrix Type 40E p H meter or a Sargent Model LS p H meter. The electrochemical cell used was a Metrohm Model EA 874 titration vessel and the indicator electrode was a Metrohm Hanging Mercury Drop Electrode ( H M D E ) , Type E410, commercially available. A micrometer dial a t the top of the H M D E assembly permits formation of mercury drops of known and reproducible size. All measurements were taken with a n electrode area of 0.0222 f 0.0007 cm2. T h e other two electrodes in the three-electrode system were a Sargent platinum electrode which served a s the counter electrode and a Sargent saturated calomel electrode (SCE) which was the reference electrode. All experiments were performed a t ambient temperature (22-23 “C). Reagents. The bilirubin used was obtained from Sigma Chemical and was of 99% purity and had a measured molar absorptivity in chloroform of 60,100 f 600. It was stored in the freezer compartment of a laboratory refrigerator. The buffer solutions prepared and used were checked against standard buffer solutions furnished by Sargent. All were 0.05M in phosphate ion. The metal ion solutions were prepared by acid dissolution of the pure metal in the cases of zinc, copper, nickel, and iron. The cobalt(I1) solutions were obtained by dissolving cobaltous chloride (CoClZ. 6HzO) in distilled water, the cadmium solution by dissolving cadmium acetate (Cd(CzH302j2.2Hz0) in distilled water, and the mercury(I1) solution by dissolving mercuric chloride (HgC12) in distilled water. Both argon and nitrogen gases were used as inert atmospheres in this study. The argon was passed over clean copper chips a t 600 “C to remove any traces of oxygen. Procedure. The following procedure was followed for obtaining the cathodic or reductive voltammograms of bilirubin. A buffer of 7.8 or 8.0 was deoxygenated by passage of argon gas through it for 1 hour. A background reductive voltammogram was obtained for the buffer by sweeping the potential from 0.0 volts us. the S C E to -2.0 volts a t different scan rates. The current (i) between the hanging mercury drop electrode ( H M D E j and the counter electrode ( P t j was measured a s a function of applied potential ( E ) between the H M D E and the reference electrode ISCE). T h e (1) R A Velapoldi, and 0. Menis, Clfn Chem.. 17, 1165 (1971) (2) J D Van Norman, Anal Chem 45, 173 (1973) (3) G . W Ewing.J. Chem Educ 46, A717 (1969)
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ANALYTICAL CHEMISTRY, VOL
46, NO. 7, JUNE 1974
resulting i-E plots were recorded on the X-Y recorder. Using the three-electrode system minimized effects due t o solution resistance. A measured amount of a bilirubin master solution (in 0.1N NaOH) was then added to the known volume of buffer, the solution again deoxygenated for 10 minutes, and more reductive scans were taken a t a variety of scan rates. The p H of the solutions were always checked after addition of the bilirubin master solution. The procedure for studying the bilirubin-metal ion complexes was similar. Again, a background was obtained on the deoxygenated buffer. Then a n aliquot of metal ion solution was added, argon bubbled again, and more reductive scans were taken. Then small known volumes of the bilirubin solution were added to give different bilirubin concentrations. Voltammograms were obtained a t each concentration of bilirubin (Note: the metal ion concentration remains constant).
RESULTS AND DISCUSSION
In voltammetry, the height of the peak is a function of concentration, scan rate, and other factors as shown below in the Randles-Sevcik equation: i, = 269 X 10’ nJ’’ AD”’ C LJ’? for reversible waves (1) or
2.98 X 10j n(cu n)’I2 AD1’? C ul’? for irreversible waves ( 2 ) where i, is the peak current in amperes, C is the concentration of electroactive species in moles per cubic centimeter, L: is the scan rate in volts per second, A is the electrode area in square centimeters, D is the diffusion coefficient in cm2/sec, n is the number of electrons transferred per molecule and a is defined as the degree of irreversibility and can be evaluated from:
i,
=
on=
0.04771
E, 2 - E,
(3)
where E,, is the peak potential and E,, 2 is the potential where the current has the value & / 2 . It should be noted t h a t Equation 1 is for a plane lineardiffusion electrode. For a spherical electrode such as the Hanging Mercury Drop Electrode (HMDE), the correct expression is
where r is the radius of the drop, 0.042 cm in this instance. For the operating conditions of this investigation such as scan rate, drop size, and diffusion coefficient, the correction to the Randles-Sevcik equation, Equation 1, is less than 5% and, for the sake of simplicity, the correction was omitted. For a more complete discussion of voltammetry, one is referred to Meites ( 4 ) . Thus, one can use Equations 1-3 to evaluate diffusion coefficients for either reversible or irreversible systems. If a metal ion and its complex reduce reversibly to the amalgam, then measurements of the shift in half wave or full wave potential (E,>2 or I?,]) to more negative potentials with an increase in ligand concentration can be applied to determine the overall stability constant. When a fairly stable complex exists in the solution and the comMeites Polarographic Techniques second ed science New York N Y 1965 p p 4 1 3 - 1 5
(4) L
revised Inter-
Table I. Diffusion Coefficients for Bilirubin. Do,crnt/sec
Scan rate, (V, rnin)
Reduction potential, V
Peak current, pA
a N n , calcd
Reversible
Irreversible
0.5 1.0 2.5 5.0
-1.37 -1.37 -1.38 -1.39
0.40 0.53 0.80 1.13
1.26 1.26 1.19 0.95
5 . 9 4 x 10-6 6.00 X 6.85 X 6.03 X 6 . 2 1 =t0 . 3 2 X 10-6
8.49 X 6.91 X 7.92 X 1 0 . 0 x 10-6
Av
8.36
+ 0.84 X
[BRI = 1 . 1 X 10-'M: supporting electrolyte, phosphate buffer, pH 7 . 8 .
h
3.-
1.0
0.5
0 1.14
1.10
1.18
1.22
1.26
- E ( V O L T S ) v s SCE
-E
(VOLTS)
VS
SCE
Figure 1. Cathodic voltammograms of bilirubin ( 1 X 10-4M) in a phosphate buffer of p H 7.8 ( A ) Scan rate of 0.5 V;min. Ep = - 1 39 V ; (6) Scan rate of 1.0 V/min, E, = -1.40 V; (C)Scan rate of 2.5 V/min, E p = 1.40 V; ( D ) Scan rate of 5 0 V,/min, E p = 1.40 V: (E)Buffer only, scan rate of 2 5 V l m i n
plexing agent is in fair excess, the following relationship holds ( 5 ) :
where and are peak potentials for the simple and complexed ion, respectively, p is the number of ligands coordinated to the metal ion (one in the case of bilirubin) and the Pp is the overall stability constant and C is the ligand (bilirubin) concentration. If we let AE be (E,,)s - (Ep)c,a plot of AE us. log C should be a straight line with a slope of 0.059 p l m and an intercept of 0.059jn log PP
.
Electrochemical reduction a t the Hanging Mercury Drop Electrode (HMDE) was accomplished for a bilirubin solution 1.0 x 10-4M a t a potential of -1.4 volts us. the SCE. The voltammograms obtained a t various scan rates are shown in Figure 1. Diffusion coefficients were calculated using Equations 1 and 2 for reversible and irreversible systems and the values obtained as shown in Table I. The bilirubin system was experimentally found to be irreversible as the reversal of the voltage scan after reduction gave no anodic wave due to the re-oxidation of the reduced species. Irreversibility could also be inferred from the spreading out of the voltammograms. Although irreversible, the reductive process was highly reproducible. The peak potential, E,, of -1.40 volts us. an SCE in the phosphate buffer of pM 7.8 compares favorably with the value reported by Tvaroha (6) for a half wave potential of -1.48 volts in 0.1M NaOH. ( 5 ) J B Headridge. "Electrochemical Techniques for Inorganic Chemists, Academic Press. London, 1969, pp 32-33 (6) 8 Tvaroha. Naturwissenschaften. 48,99 (1961) '
Figure 2. Cathodic voltammograrns of 1 X 10-4M Z n ( l l ) solution in a 0.04M phosphate buffer of pH 8 . 0 as a function of added b iIir u b i n ( A ) [Bilirubin] = 2.5 X 10-5M. E, = -1.17 V; ( D ) [Bilirubin] = 7.5 X 1 0 - 4 ~ E,, . = -1.21 v. Note: [Bilirubin] = 6.0 X
1 0 - 5 M . E, = -1.16 V; ( E ) [Bilirubin] = 5 0 X ( C ) [Bilirubin] = 6.0 X 1 0 - 5 M E,, = - 1 18 V; 1 0 - 5 M El, = -1.20 V: ( E ) [Bilirubin] = 4 2 X 10-4M. same as ( E )
Table 11. Zinc (11)-Bilirubin Complex Reduction"
[BR] moles/l.
Zn(I1) reduction potential, V
0 2 . 5 x 10-5 5 . 0 x 10-5 6 . 0 X 10-6 7 . 5 x 10-5 4 . 0 x 10-4 6 . 0 X 10-4
-1.12 -1.16 -1.17 -1.18 -1.20 -1.21 -1.215
(El>)58P
V
...
...
0.04 0.05 0.06 0.08 0.09 0,095
Av
1 . 0 x 106 1.0 x 106 2 . 0 x 106 7 . 0 X 106 3 . 0 X lo6 3 . 0 X 106 2 . 8 * 1 . 5 X lo6
[Zn(II)I = 1.0 X lO-AM, supporting electrolyte, 0 . 0 5 M phosphate buffer, pH 8 . 0 , scan rate = 2 . 6 V;min.
Diffusion coefficients were calculated for a theoretically reversible system and for the actual irreversible system. Values obtained were 6.21 & 0.32 X cm2/sec for the cm2/sec for the acreversible case and 8.36 f 0.84 x tual system. These results compare quite favorably with values found for other large organic species in aqueous solution (7). A zinc ion solution, 1.0 x 10-4M, was reduced a t the HMDE and a wave was obtained a t -1.12 volts us. the SCE. The reduction potential, E,, shifted negative as a function of added bilirubin as shown in Figure 2. Stability or formation constants for the reaction Zn(I1) + BR = ZnBR where BR represents the bilirubin species in solution and ZnBR, the complex without reference to charges, were calculated using Equation 5. The results of these cal(7) R . N Adams, "Electrochemistry at Solid Electrodes, ker, New York, N.Y , 1961, pp 220-22
Marcel Dek-
A N A L Y T I C A L C H E M I S T R Y , V O L . 46, NO. 7 , JUNE 1974
927
Table 111. Cadmium (11)-Bilirubin Complex Reductioncz
I__
[Bilirubin], moles 1.
Cd(I1) reduction potential, V
0 5 . 0 X 10-j 6 . 0 x 10-5 7.0 X 1 0 - j 4 . 0 x 10-4 5 . 0 x 10-4
-0.71 -0.73 -0.74 -0.75 -0.76 -0.76
0
0 32
0 04
... 0.02 0.03 0.04 0.05 0.05
Av
lo-
10-61
(E,,)? (El,)c,V
0.06
0.08
0.19
0.12
0. 1
PP
... 1.0x 105 1.8x 105 3 . 0 x 105
1 . 0 x 105 1 . 0 x 105
1.6
= 0.7
X lo5
[ C d ( I I ) ] = 1 . 0 X 10-’M, supporting electrolyte, 0 . 0 5 M phosphate buffer, pH 8 . 0 , scan rate, 2 5 V min. I
A E (volts)
Figure 3. Plot of [BR] vs. A€
Table IV. Iron (II)-Bilirubin Complex Reductions
culations are summarized in Table 11. This table is the result of point by point calculation. The data were also analyzed graphically. Figure 3 shows a plot of concentration of bilirubin, [BR], us. AE plotted with the ordinate on a logarithmic scale. This manner of presentation is convenient from the point of view of discussion. This plot gives a straight line with the exception of the point a t a [BR] of 1.5 x 10- 5 ~ . Equation 5 holds strictly, when the ligand concentration is in great excess. The reason for this is that when the metal ion complex is reduced to the metal a t the surface of the electrode, the ligand is released resulting in a local excess a t the electrode surface. When the ligand concentration is large compared to the metal ion, the effect of the local excess on the potential shift is small; and when the ligand is small compared to the metal ion, there is not enough to complex a significant portion of the remaining free metal ion. It is when the metal ion and ligand concentrations are similar that a maximum effect would be expected, as is the case in the present study. However, it must be recognized that the points obtained a t the lower concentrations are not as significant as the points obtained with the ligand in excess. Therefore to extrapolate to zero concentration would unfortunately magnify any errors made a t lower concentrations. For example, there is a trend noticeable in Table I1 toward lower p’s a t lower concentrations. Extrapolation of a plot of AE GS. log [BR] (not shown) yielded a Pp of 4.2 X lo5 in contrast to the average of point by point calculations of 2.8 x 106. Obviously the best solution would have been to work only at higher ligand concentrations, but unfortunately the solubility of bilirubin under these conditions is just above the highest point measured. Attempts to study zinc ion a t significantly lower concentrations resulted in the experimental inability to measure potential shifts accurately enough t o make meaningful calculations. Thus, the measurements are a compromise between the solubility limits of bilirubin and the experimental determination of potential shifts. The slope of the plot of S E GS. log [BR] gave a slope of 0.040 V which gives a value of 1.3 for “p”, interpreted as only one ligand. The trend a t higher ligand concentrations is toward a lower slope; toward p = 1. Experimental evidence for the existence of bilirubin complexes with metal ions with more than one bilirubin per metal ion is not available. There is, however, evidence that more than one metal ion may complex with bilirubin, a t least in nonaqueous media ( 8 ) . Whether this could occur in aqueous solution is a t the moment conjectural. ( 8 ) C C Kuenzle. R P Pelloni, M H Weibel, and P Hemmerich, BJOchem J 130. 1147 11972)
928
A N A L Y T I C A L C H E M I S T R Y . V O L . 46. N O
7. J U N E 1974
Fe(II1, reduction
[BR],moles ‘I.
potential, V
0 3 . 0 x 10-5 5 . 0 x 10-5 7 . 0 X 10-5 9 . 0 x 10-5 4 . 0 x lo-‘ 6 . 0 X 10-4
-0.40 -0.48 -0.49 -0.50 -0.51 -0.52 -0.53
(E,) (ED)c, V
PI,
...
,.. 0.08 0.09
2 . 0 x 10’ 2 . 2 x 107 0.10 4 . 0 x 107 0.11 6 . 0 x 107 0.12 3 . 0 X 10‘ 0.13 4 . 0 x 107 AV 3 . 5 1.1 x 107
*
[ F e ( I I ) ] = 1 . 0 x 10-1,supporting electrolyte, 0 . 0 5 M phosphate buffer, p H 8 . 0 , scan rate, 2 5 V /min.
The reduction wave of cadmium(I1) solution, 1.0 x 10-4M, occurred a t a peak potential of -0.71 volt us. the SCE. The voltammetric waves were quite similar to those of Zn(1I) except for the potential. Bilirubin-cadmium ion complex formation constants were evaluated by the shift in reduction potential and the results are shown in Table 111. The plot of AE L‘S log [BR] for the cadmium(I1) system had considerable scatter. Extrapolation to 1E = 0 of various “best straight lines” gave a range in value of Pp from 2 X lo5 to 5 X lo5 compared to the average of 1.6 X lo5 in Table 111. Slopes varied between 0.027 and 0.038 V indicating once again one ligand. As was the case for the zinc(I1) system, the point where the ligand and metal ion concentrations were similar, was well off the straight line. Iron(I1) solution, 1.0 x 10-4M, gave a reduction potential a t -0.40 volt cs. the SCE. The results for the evaluation of the formation constant of the iron(I1)-bilirubin complex are summarized in Table IV. The plot of A E us. log [BR] gave a slope of 0.033 V, indicating one ligand, and a JP of 1.25 x lo7 compared to the value in Table IV of 3.5 X 10:. Unsuccessful attempts were made to study the formation of bilirubin complexes with copper(II), mercury(II), nickel(II), and cobalt(I1). Copper(I1) reduced a t -0.46 volt but addition of bilirubin did not change the potential; however, the bilirubin yellow color changed rapidly to a green color which was attributed to the oxidation of the bilirubin to the green biliverdin. Mercury(I1) reduction wave could not be obtained with the HMDE under these conditions, and bilirubin was oxidized. Kickel(I1) waves could not be obtained either. The cobalt(I1) wave occurred a t -1.27 volts but addition of bilirubin did not shift the wave a t all. The reduction potential of the metal ions studied all were well within the range noted for each ion in a variety of solutions as reported in the literature ( 4 ) . From the
shifts in the potentials as a function of added bilirubin, the formation constants for bilirubin complexes with zinc(11), cadmium(II), and iron(I1) were evaluated. The values are; for zinc(I1)-bilirubin, /? = 2.8 f 1.5 X lo6; for cadmium(I1)-bilirubin, /3 = 1.6 f 0.7 X lo5; for iron(I1)bilirubin, /? = 3.5 f 1.1 X lo7. Since there are no reported values for these formation constants in the literature, no comparison is possible. The precision of the values for the formation constants was generally 15070,fairly typical for this type of measurement where a small error in measuring potential shifts results in a large error in the value. In each case, the value obtained for the formation constant a t approximately equal metal ion and bilirubin concentrations was significantly larger than those obtained with low bilirubin concentrations and with bilirubin in excess. This is probably due to local excesses of bilirubin a t the surface of the electrode following the metal ion reduction. This value was, however, included in the overall calculation of the average. The use of plots of LE us. log [BR] gave values of the same order of magnitude, but it is felt t h a t extrapolation
may overemphasize errors a t lower concentrations of ligand. The results obtained indicate an order of complex formation of metal ion with bilirubin of Fe(I1) > Zn(I1) > Cd(I1). It should be pointed out that the formation constants measured in this study are “effective” or “conditional” formation (stability) constants as their numerical value depends upon the solution conditions, pH, and composition. Bilirubin, however, probably does not form strong metal ion complexes in more acidic aqueous solution due to inter- and intramolecular hydrogen bonding, making measurements in acidic solution even more difficult.
ACKNOWLEDGMENT
The authors would like to acknowledge helpful discussions with L. B. Spiegel and T . N.Dobbelstein. Received for review November 21, 1973. Accepted February 27, 1974.
Induced Colorimetric Method for Carbon Monoxide Jack L. Lambert and Robert E. Wiens’ D e p a r t m e n t of Chemistry, Kansas State University, Manhattan, Kansas
Methods in common use for carbon monoxide in air include a colorimetric method which employs silver p-sulfaminobenzoate complex in alkaline solution ( I ) , infrared spectrometry at 4.67 p (2), and nondispersive infrared adsorption ( 3 ) .The method described here is the second to be reported in which a soluble colored compound is formed in aqueous solution a t room temperature; the first was a method which employed tetrachloropalladate(II), ethylenediaminetetraacetatoferrate(III), and 1.10-phenanthroline and produced red-orange tris( 1,lO-phenanthroline)iron(II) cation ( 4 , 5 ) . In the reaction described here, a reagent consisting of tetrachloropalladate(II), iodate, and leuco crystal violet (4,4’,4”-methylidynetris(N,N-dimethylaniline)produces crystal violet in proportion to the concentration of carbon monoxide and the time of reaction. At pH 3.1, the reaction between iodate and leuco crystal violet is kinetically inhibited, but the sequential reduction of palladium(I1) to palladium(0) followed by reduction of iodate by palladium(0) produces hypoiodous acid which rapidly oxidizes leuco crystal violet to crystal violet. As palladium(I1) is produced in the reduction of iodate by palladium(0). no metallic palladium appears in the solution. Iodide proPresent address, Norden Laboratories, Lincoln, Neb. 68501 ( 1 ) R . G . Smith, R. J Bryan, M Feldstein, B Levadie, F. A Miller, E. R Stephens, and N G . White (Subcommittee 4 of the Intersociety Committee), Health Lab. Sci.. 7 , 75 (January Supplement) (1970) (2) Ibid., p 78. ( 3 ) Ibid., p 81 ( 4 ) J. L . Lambert and P A Hamlin, Anal. Lett.. 4, 745 (1971) (5) J L . Lambert, R R Tschorn, and P A. Hamlln, Anal. Chem.. 44, 1529 (1972)
duced by the reduction of hypoiodous acid apparently can act as a catalyst by reacting with iodate to produce hypoiodous acid, but the catalysis reaction apparently is of minor importance in the production of crystal violet. EXPERIMENTAL All solutions were prepared from deionized xvater with the purest chemicals available from commercial sources. Prepare leuco crystal violet solution, 6.85 x 10 - 4 .Id by dissolving 0.256 gram of 4,4’,4’‘-methylidynetri~(~Y,~’v‘-dimethylani!ine) ( E a s t m a n Kodak Co.) in 200 ml of water containing 2.5 ml of’ 8
