Redox Thermodynamics and Electron Transfer Reactivity of Heme by Enthalpimetry and Voltammetry Raymond Bury‘ and Joseph Jordan” Department of Chemistry, The Pennsylvania State Unlversiv, 152 Da vey Laboratory, Unlversity Park, Pennsylvania 16802
The flrst extensive lnvestlgatlon of the redox thermochemistry of protoheme Is descrlbed. A wealth of thermodynamlc Information was obtained by judlclously combining results of potentlometrlc and thermometrlc tltratlons. Heats, Glbbs free energies, and entropies of the reduction of ferrlheme to ferroheme have been determlned In dlmethylformamlde (DMF), In dlmethylacetamlde (DMA), and In mlxtures of DMF or DMA and water. The reducing agent was chromous acetate. I n order to segregate solvent effects attrlbutable to the ferroferrlheme half-reaction, the AH, AG, and AS asslgnments have been referred to the ferrocene/ferrlclnlum couple whose energetlcs are known to be vlrtually Independent of solvent. The thermodynamlc parameters of the ferroheme-ferrlheme redox process exhlblted remarkable nonmonotonlc trends as tunctlon of solvent composltlon. The results were generally conslstent wlth concomitantly observed effects of solventsolute lnteractlons on electrode kinetics. Relevant electrochemical rate constants were evaluated by polarography and cycllc voltammetry.
and mixtures of these with water). Since ferroheme and chromous acetate were prone to air-oxidation, oxygen was eliminated by storing all chemicals in an inert atmosphere (nitrogen or argon), by deaerating solvents and solutions, and by maintaining a supernate of nitrogen or argon in all experiments. Traces of oxygen were removed from the inert gases by bubbling through a solution of vanadous sulfate. Likewise the purging gases and supernates were equilibrated by bubbling through wash bottles containing solvents of compositions identical with those used in the thermochemical and electrochemical experiments. Two types of studies were performed, viz., (I) Investigations of chemical oxidation-reduction reactions between ferriheme and chromous acetate. (11) Voltammetric studies of the ferro-ferriheme redox couple. All experiments of type I were carried out in the presence of a 0.01 M sodium salicylate-O.01M salicylic acid buffer whose pH is known to be 7.94 in DMF (13). Chromous acetate proved stable under these experimental conditions,while it decomposed rapidly (e.g., by reducing water contained in the solvent) in unbuffered solutions. In experiments of type 11,O.l M tetraethylammonium perchlorate (TEAF’) or 0.1 M perchloric acid was used as supporting electrolyte. All results reported in this paper were obtained in a temperature range of (25.00 & 0.02)
“C. Oxidation-reduction properties of octahedral complexes of iron with the equatorial ligand protoporphyrin IX are of evident biological interest because the prosthetic groups of hemoproteins (“heme”) are moieties of this type. They are implicated in crucial physiological functions ( I ) including oxygen transport (by hemoglobin), oxygen storage (by myoglobin), and cellular respiration (via cytochrome c ) . Extensive information is available in classical monographs (2, 3) on the standard and formal potentials of metalloporphyrin redox couples. Because heme tends to dimerize and polymerize in aqueous solutions ( I , 4-7) many recent electrochemical studies have been carried out in the presence of nonaqueous solvents (1, 4, 8-11) where the situation is not complicated by aggregation. Surprisingly, data on the thermochemistry of oxidationreduction reactions of heme are conspicuous by their absence. In fact, only one study on the redox thermochemistry of any iron porphyrin (viz., of hematoporphyrin) has been reported in the literature (12). In the present paper, calorimetric results are presented and discussed involving the heat of reduction of ferriheme to ferroheme with chromous acetate in the aprotic solvents N,N’-dimethylformamide (DMF) and N,”-dimethyl acetamide (DMA), and in mixtures of these with water. Concomitantly, Gibbs free energies were estimated by electrochemical methods and corresponding entropies were computed. These thermodynamic findings were complemented by evaluation of electrode kinetics via cyclic voltammetry yielding illuminating insights on mechanistically significant solvent effects.
EXPERIMENTAL General Conditions. Very dilute solutions of ferriheme, ferroheme, and chromous acetate (in a range between 0.0001 M and 0.001 M) were prepared in purified solvents (DMF, DMA, Postdoctoral Scholar (1974-75) on leave from Universit6 Pierre et Marie Curie; present address, Laboratoire d‘Electrochimie, 4 place Jussieu, 75230 Paris Cedex 05, France.
Thermometric Enthalpy Titrations. Heats of reaction were determined by a titration calorimetric method, viz., thermometric enthalpy titrations (TET). Titration curves were obtained by monitoring the change in temperature, AT, in an adiabatic cell, concomitantly with the addition of chromous acetate to ferriheme. Because solutions were extremely dilute, AT was on the order of a few millidegrees. A typical “normalized curve” is illustrated in Figure 1. Temperature increments shown on the ordinate are values corrected for changes of heat capacity due to the inevitable changes (e.g., volume increments) occurring during the titration. AT was normalized to an average heat capacity, k , effective at the mid-point of the titration. Actual heat capacities were determined by appropriate joule heating calibrations carried out in situ. Heats of reaction were computed from the ordinate increment, using Equation 1
&=kAT=-nAHFO (1) where Q denotes the number of calories evolved, n the number of moles reacted at the stoichiometric end point, A H F O denotes a formal heat of reaction which corresponded to infinite dilution [identified by the superscript “0” (zero)] of the redox reactants in a specific solvent in the presence of buffer. The crucial temperature increment, AT, (and the corresponding integral heat, Q)were assigned with the aid of the self-explanatory extrapolation shown in Figure 1, which corrects satisfactorily for extraneous effects. The most important of these were the heats of mixing as is apparent from the excess reagent line (between the “end point” and the stop of titrant addition). Actual titration curves were recorded automatically by delivering solutions of titrants at a constant rate from a motor-driven syringe-buret. Temperature was monitored with the aid of a thermistor bridge which had the effective sensitivity of a thousand-junctionthermocouple. The bridge unbalance potential (proportional to A?‘‘) was displayed on the Y-axis of a strip-chart dc millivoltmeter whose time (X)-axiswas driven by a synchronous motor. Taking into account the likewise synchronous titrant delivery (and the concentration of the titrant), the abscissa was calibrated in terms of moles of titrant added. Detailed procedures and instrumentation have been described previously (14). Entropy and Gibbs Free Energy Assignments. Equilibrium constants were evaluated from potentiometric titration curves ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
1573
+
c
where denotes a function of AE, which has been tabulated by u is the potential scan rate in volt per second, D Nicholson denotes diffusion coefficientsand the subscripts Ox and Red refer to ferri- and ferro-heme respectively. The transfer coefficient a was assumed equal to 0.5. In order to specify the formal potential at which our various k assignments were effective, the following procedure was used: (a) Polarographic half-wave potentials (E1~2eXP) in the various solvents were determined experimentally at the dropping mercury electrode. (b) Corresponding reversible half-wave potentials (El/Y)were evaluated from the relationships (18, 19)
I
W
I/
!--&
(In,
E l I Z e X=PE1/2rev+ R T In 0.886 k o , / z anF D
START
(5)
(c) Formal potentials, EFo,were estimated using the equation:
1 moles
of titrant
added
Flgure 1. Typical normalized thermometric enthalpy titration curve in DMF. Appreciable slope of excess reagent line is due to the relatively large heats of dilution in nonaqeuous solvents of low dielectric constant (e
=
Density and Viscosity Measurements and Heats of Mixing. A Cannon pycnometer (volume = 25.4842 mL a t 25 “C) and an Ubbelhode viscometer were used. These devices were calibrated with triply distilled water, utilizing the following “reference values” (20)
37)
of ferriheme with chromous acetate, which were carried out under the same conditions as the corresponding thermometric enthalpy titrations. A conventional platinum indicator electrode was used. Conditional “formal” Gibbs free energy and entropy values were computed from th8 well-known thermodynamic relationships:
4’’ = 0.99707 ~7’’ = 0.8903 cp
AGFo = - RT In K A H ~ O
=
AG,O
Heats of mixing of water and DMF (and water plus DMA) were determined by an unconventional procedure, viz., by titrating pure water into pure DMF, using the same titration calorimeter as in the thermometric enthalpy titrations described previously. Chemicals. Triply re-crystallized hemin chloride (FW 651.59), supplied by Nutritional Biochemical Corporation, Cleveland,Ohio, served as the source of ferriheme. Ferroheme was prepared from hemin chloride by controlled potential electrolysis using a vigorous stirred mercury pool cathode and a Model 61-TR Wenking Potensiostat. Chromous acetate was prepared from 0.2 M aqueous chromic chloride which contained 0.05 M sulfuric acid. That solution was percolated through a Jones reductor column (amalgamed zinc, 5% mercury) into aqueous sodium acetate, yielding a red brick chromous acetate precipitate whose composition was [ C ~ ( O A C ) ~ ] ~ This H ~ Ohydrated . dimer was washed successively with water, ethanol, and ether, pulverized, dried, and stored under argon. The presence of the water of crystallization in the compound precluded, naturally, the preparation of completely anhydrous solutions of chromous acetate in DMF or DMA. However, since the chromous solutions used in this investigation were invariably very dilute (50.001 M) the contamination amounted to 20 ppm or less. This proved tolerable and results could readily be extrapolated to the pure anhydrous solvents (see section on Results below). Solvents consisted of mixtures of N,N’-dimethylformamide (DMA, dielectric constant = 36.7) or N,N’-dimethylacetamide (DMA, dielectric constant = 37.8) and water. The aprotic solvents DMF and DMA were painstakingly purified by appropriate distillation procedures (21,22).
+ TAS~O
Voltammetry. Current-voltage curves of ferriheme and ferroheme were recorded with the aid of a Model 170 Electrochemical System, supplied by Princeton Applied Research Co., Princeton, N.J. The indicator electrodes used were the dropping mercury electrode (DME) or Kemula’s hanging drop mercury electrode (HDME) for classical polarographic measurements and cyclic voltammetry, respectively. The counter electrode was a platinum wire of relatively large area. Special care was taken to minimize uncertainties due t o liquid junction potentials. With this purpose in mind, all electrochemical experiments were duplicated using two different reference electrodes, viz., (1)A conventional aqueous saturated calomel electrode (SCE) with an agar bridge saturated with aqueous KCl. (2) Pantony’s silversilver chloride half-cell (15) in DMF saturated in potassium chloride and containing 0.8 M potassium perchlorate with a salt bridge consisting of a methyl-methyl cellulose gel impregnated with a saturated solution of tetraethylammonium perchlorate (TEAP) in DMF. Experimental potentials of the indicator electrodes DME and HDME measured against each of the two reference half cells were intercompared and found to be consistent with directly measured potential differences (60-80 mV) between the SCE and Patony’s Electrode (PE), when the two reference half cells were connected via a solvent (pure DMF and/or DMF plus water) containing 0.1 M TEAP. The potential of Pantony’s electrode vs. a normal hydrogen electrode (NHE) in DMF is known to be +0.279 V (15). Accordingly all potential assignments reported in this paper have been referred to the NHE in DMF. This must-naturally-be taken into account when comparing numerical assignments in the present communication with our preliminary data reported earlier in Ref. 8, where we referred potentials to the conventional aqueous SCE. The two sets of data are in satisfactory substantive agreement: where minor differences transpire, the values in the present communication represent the definitive assignments. Heterogeneous electron transfer rate constants (KoX = k ~ =~ KO), effective at relevant formal potentials, were evaluated (16) from anodic-cathodic potential peak separations, Up, observed on cyclic voltammograms, using the equation
RESULTS Thermodynamic Assignments. Experimental findings actually obtained by potentiometric and thermometric enthalpy titrations of ferriheme with chromous acetate yielded readily AH, AG, and A S assignments for the reaction: Fe(II1) heme t Cr(I1) = Fe(I1) heme d
(4)
1574
ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
+ Cr(II1)
(7)
Any interpretation regarding solvent effects on the ferri-ferro heme couple require, naturally, their segregation from effects due to the chromic-chromous half reaction, which is, in principle, impossible without extrathermodynamic assumptions. The following procedure was used t o circumvent this difficulty. (a) Ferricinium was titrated with Cr(I1) under the same conditions (solvent, etc.) as prevailed in Reaction 7. In this
Table 1. Thermodynamic Assignmentsa in Mixed DMF-Water Solvents Concentration of water in DMF, Cw(mol L - ' )
- AHF0(kcal mol-' )
-AGFo(kcal mol") Reaction
0 1 2 3 4 5 6 8
10 12
9
- ASF0(Cal mol-' deg-' Reaction 9
12.35 19.40 20.25 20.25 20.00 19.50 19.00 16.65 12.45 7.00
33.6 55.0 57.2 56.9 55.8 54.4 52.8 45.4 31.4 13.4
Reaction
7
8
9
7
2.49 3.25 3.57 3.70 3.87 3.94 3.99 4.06 4.13 4.17
0.15 0.25 0.32 0.41 0.50 0.65 0.72 0.94 1.05 1.18
2.34 3.00 3.20 3.29 3.37 3.29 3.27 3.12 3.08 2.99
15.50 23.00 24.20 24.65 24.95 25.00 24.80 23.65 20.90 17.40
8 3.15 3.60 3.95 4.40 4.95 5.50 5.80 7.00 8.45
10.40
a The absolute precision of the data (expressed as the standard deviation of the mean of 5 replicates) was * 2 units of the last significant figure tabulated above. Estimation of the absolute accuracy is not feasible, because comparable independent assignments are not available in the literature. The data tabulated here are the first known relevant assignments.
Table 11. Thermodynamic Assignmentsa in Mixed DMA-Water Solvents
a
Concentration of water in DMA, Cw(mo1 L - ' )
7
8
9
7
8
9
-SFo(cal mol-' deg-') Reaction 9
0 1 2 4 6 8 10 12
2.80 2.90 3.05 3.20 3.35 3.50 3.75 3.95
0.18 0.27 0.35 0.50 0.85 1.09 1.42 1.80
2.62 2.63 2.70 2.70 2.50 2.41 2.33 2.15
17.70 24.90 25.80 26.85 26.00 25.05 21.90 18.20
4.25 4.70 5.05 6.20 18.65 15.60 10.70 3.80
13.45 20.20 20.75 20.65 7.35 9.45 11.20 14.40
36.3 58.9 60.5 60.2 54.1 44.2 28.0 5.5
- AGF'(kca1 mol-')
-AH~'(kcal mol-') Reaction
Reaction
See footnote a , Table I.
+r
t
0
DMF-HpO DMA-HzO C, 2 0
C,
40
mol.
manner, effective AH, AG, and AS assignments were obtained for the process: ferricinium + ck(II) = ferrocene
+ Cr(II1)
(8)
(b) By appropriately substracting the relevant heats, Gibbs free energies and entropies of Reactions 7 and 8, thermodynamic assignments were obtained for the hypothetical redox process:
+ ferrocene = Fe(I1) heme +
mol
.f-l
Figure 3. Physical properties of DMF-water solvents as function of water content (C,)
60
p-'
Figure 2. Heat of Reaction 9 as function of water present (C,) in DMF and DMA solvents
Fe(II1) heme
Vlscoslty
ferricinium
(9)
Thermodynamic parameters of Reaction 9 are not amenable to direct experimental determination because of the closeness
of the standard potentials of the two couples and to slow kinetics, which, however, presented no difficulty in the case of Reactions 7 and 8. (c) Using Strehlow's hypothesis (23),viz., that the thermodynamic parameters of the ferrocene-ferricinium couple are virtually independent of solvent composition, i t was assumed that any solvent effects observed in Reaction 9 reflected exclusively solvent-solute interactions affecting the ferriheme-ferroheme couple. Data obtained in mixtures of water with DMF and DMA are summarized in Tables I and 11. Results in DMA and DMF exhibited similar trends which are apparent from Figure 2. Figures 3 and 4 show results of corresponding measurements of physical properties of the solvent mixtures. ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
1575
Table KII. Polarographic Diffusion Coefficients of Ferriheme ( Dox) and Ferroheme Concentration DMF of water in soll o 6 Do, l o 6 DRed lo6D o x vent mixtures, CW(mo1 L - ’ ) cmz/s cm2/s OxlDR ed cm2/s 0 4.9 5.1 1.05 *.. 1 5.2 5.2 1.00 5.2 2 4.8 5.2 1.08 5.0 3 5.0 5.2 1.04 5.0 4 4.9 4.8 0.97 5.1 5 4.9 5.1 1.05 5.0 6 4.7 4.9 1.05 5.0 8 5.0 5.5 1.10 5.3 10 5.1 5.2 1.02 5.3 12 4.9 4. I 0.95 5.5 15 5.2 4.9 0.95 5.7 18 4.9 5.0 1.02 5.8 a
Supporting electrolyte 0.1 M TEAP.
-
V
I
1
1
1
I
05
,
I
1
1
IO
mole fraction of DMF
Figure 4. Heat of mixing of DMF and water as function of mole fraction of DMF
+ e = Fe(I1) heme
(10)
determined in the present study are plotted in Figure 5 as function of solvent composition. The same results were obtained in DMF and DMA solvents in the presence of comparable amounts of water. Our DMA data agree satisfactorily with earlier work by Davis and Bynum (24). The corresponding standard potentials (where kox = k ~ =~ ko) d are also shown in the figure. In order to evaluate Eo and K O in the various solvent mixtures via Equations 4, 5, and 6, the diffusion coefficients of ferriheme and of ferroheme were required. These were calculated from polarographic diffusion currents with the aid of Koutecky’s expanded Ilkovic Equation (25). The assignments which transpired are shown in Table 111.
DISCUSSION The experimental findings are accounted for as follows. (a) The electroreactive Fe(II1) heme and Fe(I1) heme species had the axial coordination illustrated in Figure 6. The divergence of the Eo and hoplots in Figure 5 when CHzO < 5 M was due to the chloride abstraction process: X
+
H++ X(Fe heme)e?= HCl
+ X(Fe heme)X
(where X denotes DMF or DMA) occurring in the presence of HC104,which is a strong acid in DMF and DMA, while HC1 is known to be a weak acid in these solvents (26). In contradistinction, chloride remained axially coordinated (as in crystalline hemin chloride) in the presence of TEAP. When C H ~> O 10 M, both X ( F e heme)Cl and X ( F e heme)X were converted to H20(Fe h e m e ) X . This was in complete accordance with previous reports in the literature (8). 1576
DMA
l o 6D
R ~
cm2/s
0 X I D Red
...
..
5.3 5.1 5.3 5.0 5.2 5.2 5.4 5.5 5.5 5.9 6.2
*
0.98 0.98 0.94 1.02 0.96 0.96 0.98 0.96 1.00
0.97 0.94
ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
t
I
-> -E
Electrode Kinetics. Specific (“standard”) rate constants k o for the electrode reaction: Fe(II1) heme
in Various Solventsa
-1
0
@Red)
200
L------
c w mol
!-I
Figure 5. Specific (formal) electrochemical rate constants and corresponding formal potentials of the ferri-ferroheme couple in DMF in the presence of water. Supporting electrolytes are specified on each curve. C, = concentration of water present
I
I
“*
I
,
Figure 6. Axial coordination of electroreactive species in solutions prepared from hemin chloride. C, = molarity of water. X = DMF or DMA. Supporting electrolyte (0.1 M HCIO4or 0.1 M TEAP) as specified
(b) The convergence of the ko values in Figure 5 (top curves) is paralleled by striking effects in the entropies of Reaction 9, as is apparent from the minimum in Figure 7. Concomitantly the solvent properties (heat of mixing viscosity and density) also exhibit corresponding maxima or leveling features (see Figures 3 and 4). However, it is worth emphasizing that specific analytic geometry features in Figures 2 , 3 , 4 , 5 , 7 (e.g., maxima in plots of physical properties vs. leveling off of rate
Work is in progress to clarify this interesting question.
LITERATURE CITED (1) A. L. Underwood and R. W. Burnett, “Electrochemistry of Biological Compounds”, Elecfroanal. Chem., 8, 1-85 (1973). (2) J. E. Falk, “Porphyrln and Metalloporphyrins”, Elsevier, Amsterdam, 1964; 2nd ed., K. M. Smith, Ed., Elsevier, Amsterdam, 1975. (3) W. M. Clark, “OxkIation-Reductlon Potentials of Organic Systems”, William and Wilkins, Baltimore, Md., 1960. (4) D. G. Davis and D. J. Orleron, Anal. Chem., 38, 179 (1966);D. G.Davis and R. F. Martln, J . Am. Chem. SOC.,88, 1365 (1966). (5) T. M. Bednarski and J. Jordan, J . Am. Chem. SOC., 86, 5690 (1964);
89, 1552 (1967). (6) S.B. Brown and I. R. Lanske, Biocbem. J., 115,279 (1969). (7) K. M. Kadish and J. Jordan, Anal. Lett., 3, 113 (1970).
I
I L
0
!
I
I
2 0
C,
!
40
mol
1
1 6 0
(8) B. A. Feinberg, M. Gross, K. M. Kadlsh, R. S. Marano, S. J. Pace, and J. Jordan, Bloelectrochem. Bloenerg., 1, 73 (1974). (9) H. R. Gygax and J. Jordan, Discuss. Faraday Soc., 45, 227 (1968). (IO) D. W. Clark and N. S. Hush, J . Am. Chem. Soc., 87, 4238 (1965). (11) G. Peychal-Heillng and G. W. Wllson, Anal. Chem., 43, 545 (1971). (12) A. C. Censullo, J. A. Lynch, D. H. Waugh, J. Jordan, “Biochemical and Clinical Applications of Titration Calorimetry and Enthalpimetric Analysis” in “Analytical Calorimetry”, R. S. Porter and J. F. Johnson, Eds., Plenum Press, New York, N.Y., Vol. 111, 1974,pp 217-235. (13) J. Juillard and R. Loubinoux, C.R . Acad. Sci. Paris, 264, 1680 (1964). (14) J. Jordan, J. K. Grime, D. H. Waugh, C. D. Miller, H. M. Collis. and D. Lohr, Anal. Chem., 48, 427A (1976). (15) G. P. Kurnar and D. A, Pantony in “Polarography 1964,Proceedings of the Third International Congress”, Macmilian, London, 1966, p 1061. (16) R. S. Nicholson, Anal. Chem., 37, 1351 (1965). (17) R. S. Nicholson and I. Shain, Anal. Chem., 38, 706 (1964). (18) J. Koutecky, Cbem. Listy, 47, 323 (1953). (19) J. Weber and J. Koutecky, Cbem. Listy, 49, 562 (1955). (20) J. F. Swindells, C. F. Snyder, R. C. Hardy, and P. E. Golden, Natl. Bur. Stand. ( U . S . ) ,Clrc. Suppl., 440, 700 (1958). (21) H. Kojima and A. J. Bard, J . Electroanal. Chem., 63, 117 (1975). (22) J. E. Prue and P. J. Sherington, Trans. Faraday Soc., 57, 1795 (1961). (23) H. Strehlow, Z. Nekfrochem., 56, 827 (1952). (24) D. G. Davis and L. M. Bynurn, Bloelectrochem. Bioenerg., 2, 184 (1975). (25) M. Olsztajn, P. Turq, and M. Chemia, J . Cblm. Phys., 87, 217 (1970). (26) P. G. Sears, R. W. Wolford, and L. R. Dawson, J . Nectrochem. Soc..
!
8 0
I-1
Flgure 7. Variation of the entropies of Reaction 9 as function of solvent composition
constants or formal potentials) are not necessarily the same. The point we wish to make is merely the qualitative observation that changes in physical solvent properties appeared to be most drastic in certain domains of solvent composition where electrochemical thermodynamic and kinetic parameters also changed a great deal. The observed solvent property effects are indicative of changes in liquid structure which are known to be due to solvent-solvent interactions (27-29). Our findings suggest that they were matched by corresponding solvent-solute interactions reflected in electrochemical behavior. Indeed, the trends in the electrode kinetics on the one hand (Figure 5, top), and in redox thermodynamics on the other hand (Figure 5, bottom, and Figure 7), both appear to point to the involvement of the solvent-solute interactions. The kinetic effects are entirely consistent with an electron transfer path via the porphyrin ring while the entropy effects can reasonably be accounted for by the solvation of the periphery of the equatorial porphyrin ligands (8). It should be noted, however, that the findings reported in the present paper (per se!) do not necessarily preclude supplementary effects due to the involvement of axial coordination orbitals of heme-iron.
103, 633 (1956). (27) R. Paul, P. S. Guraya, and B. R. Sreenathan, Ind. J. Cbem., 1, 335 (1963); 3. .. 300 11965). (28) B. G. Cox,ATJ. Parker, and W. E. Waghorne, J . Phys. Chem., 78, 1731 (1974). (29) 0.D. Bonner and U. S. Choi, J . Phys. Cbem., 78, 1723 (1974). ~~~
RECEIVED for review June 7, 1977. Accepted June 29, 1977. Presented in part before the 4th International Conference on Chemical Thermodynamics, Montpellier, France, August 26-30, 1975. Supported by the National Science Foundation (Research Grant CHE 76-21666), the National Institutes of Health (Research Grant 5R01 HL 02342 from the National Heart, Lung, and Blood Institute), and the North Atlantic Treaty Organization (NATO Research Grant RG 794).
Assay of Phenobarbital with an Ion-Selective Electrode Gary D. Carmack and Henry Freiser“ Department of Chemistty, University of Arizona, Tucson, Arizona 8572 1
A rapid and reliable phenobarbital tablet assay method was developed based on the potentiometric sensing of the phenobarbltal anion uslng a coated-wire electrode. The results obtained are in agreement with the standard USP method.
Phenobarbital (5-ethyl-5-phenylbarbituric acid) is conventionally assayed in pharmaceutical preparations using the extractive-spectrophotometric procedure specified in the U.S.
Pharmacopeia (1). A number of other analytical methods are also available including fluorimetry (21, coulometry (31, and liquid ChOmatograPhY ( 4 ) . However, these rmd.hods generally involve the Use Of more sophisticated inStI’UmentatiOn Or more complex procedures and are perhaps best suited to the analysis of complex mixtures, e.g., human sera. The development of ion-selective electrodes based on ion association systems in this laboratory (5-7) and other (8)have demonstrated that a wide variety of simple and economical analyses are possible with these sensors. ANALYTICAL CHEMISTRY, VOL. 49,NO. 11, SEPTEMBER 1977
1577
