The Magnetic Properties of Intermediates in the Reactions of

T H E MAGNETIC PROPERTIES O F INTERMEDIATES I N T H E REACTIONS OF HEMOGLOBIN' CHARLES D. CORYELLI, LINUS PAULING,

AND

RICHARD W. DODSONs

Gates and Crellin Laboratories of chemistry, California Institute of Technology, Pasadena, California Received December 15, 1958

The theory of the oxygen equilibrium of hemoglobin proposed by Pauling (10) in 1935 treats the existent data quantitatively in terms of equilibrium constants for the stepwise oxygenation of hemoglobin, and introduces theoretical relations between successive equilibrium constants. This theory is here applied in detail to the oxygen equilibrium, and the calculated concentrations of the intermediate compounds as a function of the degree of oxygenation are presented in a figure. These calculations furnish a new point of departure for the examination of physicochemical studies of the hemoglobin-oxyhemoglobin system and of systems similar in nature, such as the hemoglobin-carbonmonoxyhemoglobin system and the hemoglobin-nitric oxide hemoglobin system. The experimental portion of this investigation consists of titration studies of two systems, involving complex formation of hemoglobin with oxygen and with nitric oxide, for which the magnetic properties have been followed as a function of the degree of completion of the reaction. These measurements show that the values of the magnetic susceptibility of the intermediate compounds formed during these reactions are linearly related to the number of hemes that have undergone reaction. The observations make it possible to reject one structural explanation of the magnetic moment of ferrohemoglobin, namely, that each iron atom has two unpaired electrons and that the spin moments of the four iron atoms of the molecule are brought into alignment by chemical forces operating between the hemes; and they thereby strengthen the conclusion reached by Pauling and Coryell (12) that the iron atoms have four unpaired electrons each and are held in the hemoglobin molecule by bonds that are essentially ionic. 1 Contribution No. 665 from the Gates and Crellin Laboratories of Chemistry of the California Institute of Technology. Present address: Department of Chemistry, University of California a t Los Angeles, Los Angeles, California. * Present address: Department of Chemistry, The Johns Hopkins University, Baltimore, Maryland.

825

826

C. D. CORYELL, L. PAULING AND R. W. DODSON

This investigation is a part of the program of study of the structure of hemoglobin conducted in these Laboratories with the aid of a grant from the Rockefeller Foundation. I. INTERMEDIATES I N HEMOGLOBIN COMPOUND FORMATION

Physicochemical studies of compound formation of ferrohemoglobin with oxygen and carbon monoxide have shown that the reaction does not correspond to a simple equilibrium between each individual heme group (Hb) and the reacting substance. It is now known that under ordinary conditions in aqueous solution there are four iron-porphyrin complexes in a hemoglobin molecule (Hb4), with molecular weight close to 67,000. Adair (1) and Ferry and Green (7) have shown that the oxygen equilibrium of purified hemoglobin solutions can be interpreted quantitatively by taking into account the existence of the intermediate oxygen compounds Hb402,HbaO4, and Hb&, as well as Hb4 and Hb&, and they have determined equilibrium constants for each step of the reaction, obtaining an equation which represents empirically the overall oxygen equilibrium. Equilibrium measurements on defibrinated whole blood and on corpuscular suspensions have similarly been explained satisfactorily by the postulate of the existence of intermediate compounds. These measurements indicate the same molecular complexity for hemoglobin in its natural state in the erythrocytes as in solution with various buffers. It is to be noted that convincing evidence for the existence of intermediate compounds in hemoglobin reactions is provided by the studies by a number of observers of equilibria of hemoglobins of various species, in whole blood and in solutions of various concentrations and a t various p H values, with oxygen and carbon monoxide; and that the theory of intermediate compounds gives the only explanation of the experimental data that is compatible with the degree of aggregation (Hb4) of hemoglobin. Pauling in 1935 (10) analyzed the extensive measurements of Ferry and Green (7) on horse hemoglobin and showed that the results are compatible with the assumption of interactions operative between the heme groups in such a manner that the free energy of oxygenation of one heme group is decreased by RT In a when an adjacent heme group is already combined with an oxygen molecule. Two spatial arrangements of the hemes were considered in detail,-(a) a t the corners of a square and ( b ) a t the corners of a tetrahedron. The latter possibility was rejected, because it would require the heme groups to be some 47 A. apart on the surface of the globin molecule and because it would require each heme to be bonded to three others, which is rendered improbable by the known structure of the porphyrin molecule. From the dependence of the oxygen equilibrium constant of the heme on acidity it was also concluded that acid groups whose ionization constants are increased by oxygenation are coupled with

827

INTERMEDIATES I N REACTIONS O F HEMOGLOBIN

the heme groups. The total heme-heme interaction energy in the molecule was found to be about 6000 cal. per mole, and the total heme-acid group interaction energy about 6600 cal. per mole. Assuming the square arrangement of hemes, the heme-heme interaction energy corresponds to the value 12 for the factor a,which represents the increase in the equilibrium constant of the heme-oxygen reaction when an adjacent heme is already in combination. We shall use the name interaction constant for a. Many experimenters have tried to find direct evidence for the existence of the intermediate compounds in reactions such as that of oxygenation. (See, for instance, the paper of Conant and McGrew (4).) The interest in these investigations and the importance of outlining the precision with which experiments must be made in order to test directly the interaction theory make it desirable to calculate the concentrations of the various intermediates in the oxygenation of hemoglobin a t various stages of the reaction, as predicted by this theory. The equations used in the calculations will be derived below. It is convenient to make the calculations for horse hemoglobin, assuming square configuration of the heme groups4 and the experimental value 12 for a as determined from the data of Ferry and Green on horse hemoglobin. It is not unlikely that the numerical value of the interaction constant a will be found to vary somewhat with the nature of the globin, since the interaction system probably involves the protein molecule as well as the prosthetic group. The relative concentrations of the two end substances and the various intermediates are given a t constant acidity by the following values, with the concentration co of Hbr taken as standard a t unity: Hb-Hb 1 1 Hb-Hb

Hb-H bO2

I

Hb---Hb

co = 1

CI =

Hb-Hb02

I

OzHb-Hb

1

1

4Kp

Hb-HbOz 1

1

OzHb-H b 0 2 cIII =

Hb-HbOz Hi, -Hi CII =

o2

4aK2p2

(1)

OzHb-HbOz

1

'

OzHb-H b 02

4a2K3p3

4 We have also made calculations of the concentrations of the various intermediates in the oxygenation of hemoglobin, assuming the tetrahedral configuration of the hemes. These calculations give nearly the same results as those given below for the concentrations of Hbd, HbrOt, HbdOs, and HbrOs, but the concentration of the dioxy intermediate H b 4 0 d is approximately two-thirds as great t h r o u g h n t the whole range. We know of no measurements sufficiently precise t o distinguish between these two possible configurations, and we shall restrict ourselves to consideration of the square configuration as the more probable one and a s a close approximation t o the truth unless i e w evidence conflicting with this is brought to light.

828

C. D. CORYELL, L. PAULINQ AND R . W . DODSON

The numbers 4, 4 , 2, and 4 are the statistical weights of the intermediates (the numbers of ways in which the oxygen molecules can be added to the hemoglobin molecule to give the intermediates). The symbol p represents the oxygen pressure, in millimeters of mercury, and K the equilibrium constant for oxygenation of an isolated heme, as defined by the equation

Its value is dependent on the acidity of the solution (10). The relative concentration of an intermediate is increased a-fold for each pair of adjacent hemes contained in the molecule, independent of the acidity. With consideration of the oxygen content of each molecular species, the values for the relative concentrations given in equation 1 lead to the following equation for the fraction of saturation, y, as a function of the oxygen pressure : Kp (2a 1) K z p 2 3a2K3ps a 4 K 4 p 4 (3) = 1 4Kp + (4a 2 ) Kzpz 4aZK8ps a4K4p4

+

+

+

+

+

+

+

+

This equation, with the value 12 for a , represents very well the observations of Ferry and Green. From equations 1 and 3 there can be derived expressions for the concentrations of the molecular species as functions of the fractional saturation with oxygen. The concentration of a given species for a given value of K p is divided by the sum of all concentrations to give a normalized value; curves representing these quantities are plotted against y in figure 1. The values of the ordinates in figure 1 give, accordingly, the fraction of the hemoglobin molecules existing in a given oxygenation step for the corresponding values of the overall saturation plotted as abscissas. This theory of the effect of hemoglobin structure on physicochemical relationships accounts for the characteristic sigmoid saturation curves of hemoglobin solutions and whole blood with oxygen and carbon monoxide in the following manner: At very low oxygen pressures the principal oxygen compound must be the monoxy compoun?. The formation of the cis-dioxy compound is made easier than would be expected from statistical calculations for independent hemes, because it contains a pair of adjacent oxyhemes which stabilizes the molecule by the amount of free energy RT In a,and consequently the hemoglobin solution takes up oxygen more readily after an appreciable amount of monoxy compound has been formed than a t the start of the reaction. This causes the upturn of the oxygen saturation curve in the early stages of the reaction. The trioxy compound is stabilized over the cis-dioxy compound by the same amount that the cis-dioxy compound is stabilized over the monoxy compound, whereas the tetroxy compound is doubly stabilized over the trioxy compound because


829

it has four interacting pairs of hemes instead of two; this compound becomes of importance even early in the reaction, and its stability accounts in the main for the rapid approach of the sigmoid saturation curve to the asymptotic value. It is seen that at half-saturation 34 per cent of the hemoglobin is in the form of intermediates. The saturation curve of carbon monoxide with hemoglobin is also sigmoid; and the relatively few experiments that have been reported indicate that the degree of sigmoid character is very close t o or identical with that for the oxygen curve (see the curves of Douglas, Haldane, and Haldane (6) for oxygen and for carbon monoxide). Further qualitative evidence for the close similarity in degree of sigmoid character of the two curves is

‘i

FIG.1. Concentrations of molecular species as functions of degree of oxygenation of horse hemoglobin. 0 = H b r ; I = HbdOg; I1 = cis-Hbr0g; I11 = HbdOs; IV = HbrOs. The curve for trans-Hb4Oris not given; it has one twenty-fourth the height of I1 throughout.

found in the study of the reaction between oxyhemoglobin and carbon monoxide (6), in which the hemes appear to act independently. If a for the carbon monoxidcl equilibrium of horse hemoglobin is equal to that for the oxygen equilibrium, figure 1 is directly applicable to the distribution of molecular species in the carbon monoxide equilibrium; it is in any case applicable as a fair approximation. Since nitric oxide hemoglobin has structure and physical properties nearly like those of oxyhemoglobin and carbonmonoxyhemoglobin, it is probable that its equilibrium with hemoglobin also can be represented approximately, if not exactly, by equation 3 and figure I . Figure 1 may be used in general in correlating with the degree of reaction properties of hemoglobin which are affected by reaction with complex-

830


forming molecules. These properties may include magnetic susceptibility, base-binding capacity, chemical reactivity, heat of formation, and others. 11. MAGNETIC TITRATIONS INVOLVING FERROHEMOGLOBIN

The value of the magnetic susceptibility of cow hemoglobin a t 25OC. has been redetermined by Taylor and Coryell (13) and found to be 4 X (12,290 f 60) X It was found that the magnetic susceptibilities of horse and sheep hemoglobins have the same value, to within the experimental errors of their determination, whereas that of human hemoglobin is about 3 per cent lower. Pauling and Coryell (12) discussed in a structural interpretation of the magnetic properties of ferrohemoglobin compounds two alternative possibilities, I and 11,in explanation of this value. I t was assumed as hypothesis I that the chemical interactions between the ferrohemes are strong enough to couple the magnetic moments of the four iron atoms into a resultant moment, and that each iron atom contributes two unpaired electrons, leading to a magnetic moment for the hemoglobin molecule of 8.95 Bohr magnetons plus about 10 per cent orbital contribution,6 or, roughly, 10 magnetons. This value is to be compared with the value 10.87 magnetons calculated from the experimental value of the susceptibility of the hemoglobin molecule, assuming the conditions of hypothesis I. As the alternative hypothesis 11, it was assumed that the iron atoms are independent in their magnetic behavior, leading to a predicted moment per heme group of 4.90 plus orbital contribution. This is to be compared with the value 5.44 calculated from the experiments assuming the conditions of the hypothesis. (Observed values for ionic ferrous compounds are about 5.1 to 5.3 magnetons.) Hypothesis I1 was accepted as the more probable one by these workers, partly because of the results of experiments preliminary to those presented in this investigation. We shall explore the first possibility further. We assume that each ferroheme group contributes two unpaired electrons to the total number, and that unpaired electrons are coupled with parallel moment vectors for ferroheme groups. We expect then for a series of cumpounds involving 6 The magnetic moments of the iron group elements are due principally t o the spin moments of unpaired electrons. The total spin moment, p', is given by

Po =

d n ( n 4- 2)

where IE is the number of unpaired electrons. The orbital contribution to the magnetic moment is nearly completely quenched, but enough remains to raise the value of the magnetic moment somewhat (for elements with more than five electrons in the 3d levels), by amounts that as yet cannot be predicted accurately, since they depend on chemical environment as well as on the spectroscopic term value of the paramagnetic atom.

INTERMEDIATES IN REACTIONS OF HEMOGLOBIN

831

transformation of hemes from paramagnetic to diamagnetic, such as for the series Hb4, Hb402,HbaOl, HbnOB, and HbaOs, spin moment values of 1/48,1/24, 1//sj and 0. We make the assumption, reasonable for such a series of closely related compounds, that the orbital moment persisting&is a constant fraction of the total spin moment. The value of the molal paramagnetic susceptibility for each of these compounds, which is directly proportional to the force determined by the Gouy method, is proportional to the square of the moments (given above) for them. Since measured magnetic susceptibilities are proportional to concentration and are additive, the magnetic susceptibility of a hemoglobin solution

da

Y

FIG. 2

ML

N.,S20+ SOLUTDN

FIG.3

FIG.2. Predicted magnetic susceptibility of a hemoglobin solution as a function of degree of saturation with oxygen. FIG.3. Magnetic titration of oxyhemoglobin solution with sodium hydrosulfite solution. Run No. 1.

as a function of the degree of saturation, y, with a compound giving diamagnetic iron atoms can be predicted from these values for molal susceptibilities and the values for concentrations of various hemoglobin derivatives given in figure 1. A plot of the predicted dependence of the susceptibility kx on y is given in figure 2.6 The solid curve represents the predictions for hypothesis I with complete magnetic interactions, and the dashed line that for complete magnetic independence of the hemes. It is seen that a t y = 0.5 the magnetic susceptibility would be 11 per cent e The assumption of tetrahedtal heme configuration (see footnote 4) would lead to a curve nearly identical with this one.

832

c. D. COBYELL, L. ~ A U ~ I NAND G A. w. DODSON

lower than that expected for linear dependence of magnetic susceptibility on y. I n applying figure 2 to the susceptibility measurements made on cow hemoglobin we make the assumption that a has the value 12 for the oxygen equilibrium of this substance; this assumption is a fair one, for measurements of Brown and Hill (3) show that the degree of sigmoid character of the oxygen saturation curve for cow hemoglobin is approximately the same as that for horse hemoglobin. Titrations of oxyhemoglobin with hydrosulfite Solutions of oxyhemoglobin were titrated a t a room temperature of about 26OC. with standard solutions of sodium hydrosulfite (Na&04). The magnetic susceptibilities for each point during the course of the reduction of oxyhemoglobin to ferrohemoglobin were measured by the Gouy method. (See earlier papers by the authors for descriptions of the method of preparation of cow oxyhemoglobin solutions (12) and of determination of the magnetic susceptibility (11) .) The sodium hydrosulfite solutions were prepared by dissolving about 3 g. of the salt of 85 per cent purity (from the Eastman Kodak Company) and 1.5 g. of anhydrous sodium carbonate in 20 ml. of distilled water. These solutions were kept in a test tube under a thin rubber stopper, and were freshly prepared for use. Small portions of the hydrosulfite were withdrawn in a medical syringe provided with a stainless-steel hollow needle and graduated to 0.01 ml., and were injected through a sliding rubber piston into about 33 ml. of the oxyhemoglobin solution, which was kept without gas phase in the differential susceptibility tube. The tube wtw rotated end-over-end by motor for 5 min. after each addition of reagent, with glass beads present to aid in stirring. The concentration of the sodium hydrosulfite solution was checked a t the beginning and end of each titration by titrating against a pipetted volume of potassium triiodide solution, also under a rubber piston, to the disappearance of the iodine color. From 0.5 to 0.8 ml. of the sodium hydrosulfite solution reduced 20 ml. of the 0.1 N triiodide solution. One oxyhemoglobin titration took from 4 to 7 hr. Forces in milligrams for a standard magnetic field ( Aw) have to be corrected by small quantities for diamagnetism of the reagent added (experimentally determined) and for dilution of the hemoglobin solution, in order to be proportional to the magnetic susceptibilities. There are presented in figure 3 the corrected values of Aw of run No. 1 plotted against the volume of hydrosulfite solution added. A straight line has been drawn through the points taken during the reduction reaction, and this has been continued as a horizontal line after the end point taken a t 1.47 ml. Consideration8 of apparatus and technique lead us to believe that


833

the error in Aw may amount to 0.1 and the error in volume in any addition of hydrosulfite may amount to =k 0.005 ml. The value of Aw of the initial point differs from zero because the oxyhemoglobin solution is more diamagnetic than water, the experimental reference substance. The average paramagnetic susceptibility of the iron in the solution is proportional to the increase in the corrected Aw over the initial valurl. At the beginning of the titration a n average of 0.79 ml. of hydrosulfite reduced 20.0 ml. of standard triiodide, and a t the end of the titration an average of 0.80 ml. was necessary. It is calculated from these titrations that the effect of decomposition of hydrosulfite on the curve is negligible. Another source of error is the transformation of oxyhemoglobin to ferrihemoglobin at low oxygen pressures (2). Experiments were carried out which showed that this effect caused a n increase in Aw of about 0.2 per hour for the middle portion of the curve under the experimental conditions here prevailing. This effect would operate to make the experimental curve slightly concave upwards, but the curvature mould be so small as to escape notice. We conclude, therefore, that transformation of oxyhemoglobin to ferrihemoglobin is not a yerious source of error in these experiments. Another oxyhemoglobin solution 0.0133 f in heme iron (concentration determined magnetically) was titrated in the qame manner, the titration being repeated the next day. The data, corrected for diamagnetism of reagent and dilution, are presented in figure 4. I n the first titration, run KO.2, the amounts of hydrosulfite required to reduce 20.0 ml. of triiodide were initially 0.47 and finally 0 50 ml. I n run X o 3 the respective titers of the same hydrosulfite solution were 0.53 and 0.57 ml.. The magnetic end points came a t 0.85 and 1.18 ml. for the two runs, as seen in figure 4. Before these runs can be interpreted, corrections must be applied for the large change in concentration of hydrosulfite which occurred during the titrations. Meyer (8) haq shonn that 1 mole of hydrosulfite reduces nearly 1 mole of oxygen to give principally 1 mole of sulfate and 1 mole of sulfite, according to the following equation.

+

+

+

+

02 S204-H2O = SO3-SO,-2Hf (5) Experiments of Xicloux and Roche (9) include a titration of dog oxyhemoglobin with hydrosulfite with results u hich correspond to this equation much more closely than do those of Meyer, the better results probably being due to the buffering action of the blood proteins. We can calculate then from a n oxyhemoglobin titration the average concentration of hydrosulfite in the reagent.

834


Meyer also showed (8) that 2 moles of hydrosulfite on decomposition in solution yield 1 mole of thiosulfate and 2 moles of sulfite, 2&04--

+ HZ0 = &Oa-- + 2509-- + 2H+

(6)

and that the decomposition is catalyzed by the decomposition products. Hydrosulfite, thiosulfate, and sulfite ions reduce respectively 6, 1, and 2 equivalents of triiodide per mole. We calculate from the above information that 3.5 equivalents of reducing power to triiodide disappear per mole of hydrosulfite decomposed. Using the oxyhemoglobin titer and the initial and final triiodide titers, we have calculated accordingly

ML. N&O+

FIG.4

SOLUTW

ML. N&O,

SOLUTION. CORRECTED

FIG. 5

FIG.4. Magnetic titrations of oxyhemoglobin solution with sodium hydrosulfite solution. 0, run No. 2; 0, run No. 3. FIG.5 . Magnetic titrations of oxyhemoglobin solution with sodium hydrosulfite solution, with correction for the decomposition of the hydrosulfite. Q , run N o . 2 ; 0, run No. 3, with abscissas changed. that the hydrosulfite concentration decreased 14 per cent during run No. 2 and 19 per cent during run KO.3. The period of time between initial and final titration of the triiodide solution was for each run 53 hr. Assuming that the rate of decomposition of the solution was essentially constant, we have calculated the concentration of the hydrosulfite solution at the recorded time of each addition, and have expressed the amount added in terms of the volume of the hydrosulfite solution a t the concentration prevailing a t the start of each run. The end points for the two runs with volume of hydrosulfite expressed in this manner are a t 0.79 ml. and 1.10 rnl., respectively. We have multiplied the volumes for run No. 3 by the


835

factor 0.79/1.10 to make them directly comparable with the results of run No. 2, and have presented the runs together in figure 5, with a straight line connecting the initial point and the magnetic end point for comparison with the experimental results. In a fourth run with a third oxyhemoglobin solution Aw values lie very accurately on a straight line when plotted against volume of hydrosulfite up to three-quarters reduction; during this time very little decomposition of hydrosulfite occurred. At this point the hemoglobin solution was left overnight in the refrigerator, and the titration was continued the next

ML NINO,

SOLUTION

FIG.7 FIG.6. Magnetic titration of oxyhemoglobin solution with sodium hydrosulfite solution. Run No. 4. The dashed portion of the curve was obtained the second day. FIG.7. Magnetic titrations of hemoglobin solution with sodium nitrite solution at pH 5.5 in the presence of excess hydrosulfite. 0, run No. 11; 8 , run No. 12; 6 ,run No. 13; C), run No. 14.

morning, a t which time ferrihemoglobin formation had increased Aw by 0.50 and the hydrosulfite solution had become more dilute. A quantitative determination of the end point in terms of initial oxyhemoglobin and hydrosulfite concentrations cannot be made, but the portion of this run completed the first day gave better linear dependence than any of the three other runs reported, and the extrapolated end point (1.40 ml.) agrtes within 0.05 ml. with the end point obtained the next day in the presence of ferrihemoglobin with the then more dilute hydrosulfite solution. The results of this run are presented in figure 6, with broken circles representing the unreliable data of the second day. The straight line representing points of the first day is extrapolated to give the end point.

836

C. D. CORYELL, L. PAULING AND R. W. DODBON

Conclusions about the oxyhemoglobin-hydrosulfite titrations will be discussed in the third section of this paper. Titrations of ferrohemoglobin with nitric oxide Solutions of ferrohemoglobin in the presence of hydrosulfite were titrated with a standard solution of sodium nitrite to produce nitric oxide hemoglobin. Ferrihemoglobin solution, used as the source of ferrohemoglobin, was reduced with 0.6-0.9 g. of sodium hydrosulfite. The hemoglobin was retained under a rubber piston, with absence of gas phase, as in the previous experiments with oxyhemoglobin. The concentration of the sodium nitrite solution was determined by oxidation in hot solution with excess standard permanganate and acid, followed by titration with thiosulfate after the addition of iodide. Three titrations gave a mean concentration of 0.353 f. Meyer (8) found that alkaline hydrosulfite solution docs not reduce nitrite and that an excess of hydrosulfite in acid solutions gives nitrous oxide or nitrogen. Nitric oxide is produced in very acid solutions or with an excess of nitrite. Experiments of Coryell and Dodson (5) show, however, that nitric oxide ferrohemoglobin readily forms when hemoglobin is present with hydrosulfite and nitrite, probably because of the high affinity of ferrohemoglobin for nitric oxide and the extreme stability of the compound with respect to reduction. It was also shown that nitric oxide hemoglobin is closely similar in structure to carbon monoxyhemoglobin and oxyhemoglobin in that the iron atoms are involved in octahedral covalent bond formation. There remains, however, one unpaired electron per heme, owing to the odd number of electrons in the NO ;roup. The molal susceptibility of nitric oxide hemoglobin (per heme) was found to be 1280 x 10-6, that is, about one-tenth of the value for ferrohemoglobin. Attempts were made to define conditions under which nitrite would be reduced by hydrosulfite in the presence of hemoglobin to give stoichiometric quantities of nitric oxide hemoglobin. Additions of nitrite to hemoglobin-hydrosulfite mixture with bicarbonate-carbonate buffer a t pH 9.5 led to only a small reduction in susceptibility even over a period of several days, the average fall in Aw being only 7 per cent of that expected for stoichiometric formation of nitric oxide hemoglobin. Addition of nitrite to hemoglobin-hydrosulfite mixture with acetic acid-acetate buffer a t pH of approximately 4.7 led to a large fall of Aw, corresponding to nitric oxide hemoglobin formation, followed by a rise on addition of a second portion. The spectrum then showed the presence of ferrihemoglobin as well as nitric oxide hemoglobin. On longer standing the susceptibility fell and considerable denaturation occurred. It was concluded that titrations at this low pH value could not be carried out because of the rapid destruction of hydrosulfite, leading to slow reduction of nitrite

INTERMEDIATE& I N REACTIONS OF HEMOGLOBIN

837

by other reducing agents, including hemoglobin, and because of the occurrence of denaturation. It was found that titrations could be made smoothly starting with 35.0 ml. of ferrihemoglobin 0.0147 f in heme iron and reduced with a t least 0.9 g. of sodium hydrosulfite. The pH of the ferrihemoglobin stock solution was 6.35 (by glass electrode), and the pH during the titration was near 5.5. The titration did not proceed to an end point involving complete formation of nitric oxide hemoglobin when only 0.6 g. of hydrosulfite had been added, but the addition of more hydrosulfite to such solutions with enough nitrite yielded pure nitric oxide hemoglobin (runs 11 and 12). The results of four such runs (11 to 14) are presented in figure 7. Runs 13 and 14 are considered to be more accurate than the others, because the volume of nitrite was computed from the increase in weight of the susceptibility tube and the density of the solution, 1.014 a t 25"C., and because they extend over the whole range. Values of Aw have been corrected for dilution of the hemoglobin solution only, since the diamagnetism of the nitrite solution was found not to differ appreciably from that of water. The end point for the nitrite runs was found to be a t 1.595 f 0.005 ml. for 35.0 mi. of hemoglobin. The end point predicted for stoichiometric reduction of nitrite to nitric oxide is calculated to be at 1.46 ml. This discrepancy of 9.2 per cent arises from reduction of nitrite in part to some substance other than nitric oxide, probably nitrous oxide (8). It is very probable that the fraction of nitrite involved in the subsidiary reaction remains constant during a run. 111. DISCUSSION

The magnetic titrations involving the oxygen-hemoglobin reaction are in definite disagreement with hypothesis I for the structure of ferrohemoglobin, which predicts a pronounced upward concavity of the curve of susceptibility plotted against degree of completion of the reaction. The data presented in figures 3 and 6 show the dependence of susceptibility on degree of oxygenation to be linear to within the experimental error of measurement. The results of figure 5 do not exclude a small upward concavity, but because of the large corrections for progressive change in hydrosulfite concentration these data are much less reliable. The application of the curve of figure 2 as representative of the predictions of hypothesis I for the nitric oxide hemoglobin reaction cannot be made so directly. The magnetic susceptibility of nitric oxide ferrohemoglobin (5) is incompatible with the existence of complete magnetic interactions for this substance, for this postulate (the analog of hypothesis I for ferrohemoglobin) would require a fraction of a free electron per heme, whereas complete agreement is obtained with the assumption of independent hemes with one free electron apiece. We may, however, proceed

838

C. D. CORYELL, L. PAULING AND R . W.

DODSON

to test hypothesis I for ferrohemoglobin by making the assumptions that there is complete interaction between ferrohemes but that the nitric oxide ferrohemes are independent of each other and of the ferrohemes, and that the distribution of molecular species is the same for the reaction of hemoglobin with nitric oxide as for that with oxygen (figure 1). The expected susceptibility curve would then be identical with the solid one in figure 2, except for an additional paramagnetism proportional to y to take into account the paramagnetism of the nitric oxide ferrohemes, and would show the same upward concavity' The experimental data given in figure 7 provide the best evidence yet obtained for the linear dependence of susceptibility on degree of reaction. We accordingly reject hypothesis I for ferrohemoglobin on the basis of this evidence also. It was assumed in hypothesis I1 that the ferroheme groups are essentially independent magnetically. It was, however, concluded (12) from consideration of the absolute magnetic moment of the ferrous atom, 5.435, which is higher than expected for ferrous ion by several tenths of a Bohr magneton, that there might be a small amount of magnetic interaction between the ferrohemes associated with the chemical interaction between them. The magnetic titrations. particularly those of the nitrite runs 11 to 14 shown in figure 7, indicate that this interaction has a t most only a small effect, for the amount of interaction moment would fall off rapidly as hemes are transformed from the paramagnetic to the diamagnetic state, and upward concavity would be expected for the susceptibility-degree of reaction curve in the manner predicted for hypothesis I, except that the concavity might be much less. We conclude that magnetic interactions between the hemes are not directly detectable a t room temperature by titration. Other experimental investigations of possible magnetic interaction effects are under way in these Laboratories. It has been emphasized previously (12) that the oxygenation of hemoglobin, as well as the reactions with carbon monoxide and with nitric oxide, involve a considerable change in structure, which is reflected in a pronounced change in magnetic properties. The linear dependence of magnetic susceptibility on degree of reaction with substances forming complexes shows that the structure of one heme group (ionic or covalent) does not affect that of another, so far as magnetic properties and bond type are concerned. The bond type of a ferroheme group (Hb) of a hemoglobin molecule is accordingly the same if it is part of the molecule Hb4 or of a molecule as highly oxygenated as HbaOs, and the bond type of an oxyheme group (HbOz) is the same in such diverse molecules as Hb40sand HbdOOs. IV. SUMMARY

The theory of chemical interactions between adjacent hemes in hemoglobin hcts been presented in detail, and the concentrations of various inter-

INTERMEDIATES I N REACTIONS OF HEMOGLOBIN

839

mediates in the reaction of hemoglobin with oxygen (and probably with carbon monoxide and with nitric oxide) have been calculated as a function of the degree of reaction. Magnetic measurements have been carried out during titrations involving the transformation of oxyhemoglobin to ferrohemoglobin and of ferrohemoglobin to nitric oxide hemoglobin. From the results of these it is concluded that each of the four hemes of the ferrohemoglobin molecule contains four unpaired electrons (complete ionic structure) and that the magnetic moments of the hemes are essentially independent, the alternative hypothesis of two unpaired electrons per heme with complete magnetic interaction (partially covalent structure) being rejected as incompatible with the experimental data. It is also shown that magnetic titrations a t room temperature fail to give evidence of any magnetic effect of chemical interactions between ferrohemes. The magnetic susceptibilities of intermediate compounds formed in hemoglobin reactions are accordingly linearly related to the number of hemes that have undergone reaction. A quantitative investigation of the reduction of nitrite by hydrosulfite in the presence of hemoglobin to give nitric oxide hemoglobin has been carried out. REFERENCES

(1) ADAIR,G. A.: J. Biol. Chem. 83, 529 (1925); Proc. Roy. SOC.(London) AlOQ, 299 (1925). (2) BROOKS, J.: Proc. Roy. SOC.(London) B118, 560 (1935). (3) BROWN, W.E. L., AND HILL,A. V.: Proc. Roy. SOC.(London) B94,297 (1923). (4) CONANT, J. B., AND MCGREW,R. V.: J. Biol. Chem. 86,421 (1930). (5) CORYELL, C.D., AND DODSON,R. W.: J. Biol. Chem., in press. (6) DOUGLAS, C. G., HALDANE, J. B., ANDHALDANE, J. S.: J. Physiol. 44,275 (1913). (7) FERRY,R. M., AND GREEN,A. A.: J. Biol. Chem. 81, 175 (1929). (8) MEYER,J.: 2. anorg. Chem. 34, 43 (1903). (9) NICLOUX, M.,AND ROCHE,J.: Bull. soc. chim. biol. 8, 71 (1926). (10) PAULING, L.:Proc. Nittl. Acad. Sci. U. S. 21, 186 (1935). (11) PAIJLING, L.,AND CORYELL, C. D.: Proc. Natl. Acad. Sci. U. S. 22, 159 (1936). (12) PAIJLING,L., AND CORYELL, C. D.: Proc. Natl. Acad. Sci. U. S. 22, 210 (1936). (13) TAYLOR, D.S.,AND CORYELL, C. D:. J. Am. Chem. SOC.60,1177 (1938).

The Magnetic Properties of Intermediates in the Reactions of

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