Set the tubes in a rack for color development in a water bath at 37' c for 15 minutes. Alternatively Set at room temperature (20" C or more) for 30 minutes. Measure absorbance at 625 mp. Subtract absorbance of blank and calculate concentration from absorbance of standard. ACKNOWLEDGMENT
The author thanks Margot Pickard and Thomas Chung for technical assistance in carrying out this work, and R. H.
Allen and J. E. Logan for helpful suggestions in preparation ofthe paper and for interest in the work. RECEIVED for review January 9,1967. Accepted April 3, 1967. Part of this work was presented as a paper at the Conference of the Canadian Society of Clinical Chemists, Winnipeg, Manitoba, June 1966. Work conducted in The Medical Research Section, Ottawa Civic Hospital, Ottawa, Canada.
Ion Pair Extraction of Pharmaceutical Amines Role of Dipolar Solvating Agents in Extraction of Dextromethorphan Takeru Higuchi, Arthur Michaelis, T. Tan, and Arthur Hurwitz Laboratory f o r Pharmaceutical Analysis, School of Pharmacy, University of Wisconsin, Madison, Wis. 53706 Although the ready transfer of ion pairs formed between several simple anionic species and nitrogenous cations from aqueous solution to lipoidal solvent is well recognized, the special solvating role of protondonating molecules has not been seriously considered. Since the phenomenon affords a particularly useful technique in separation and analysis of amines, the physical chemistry of these systems has been studied. Data are presented which suggest that ion pairs such as dextromethorphan hydrochloride can be extracted into organic solvents only as complexes containing of the order of five molecules of proton donors such as chloroform or phenol. The nitrate, bromide, and iodide behave similarly, except that the latter appears to require only four molecules of the proton donor. Less hydrophilic anions yielded more readily extractable ion pairs.
ALTHOUGH EXTRACTION OF PHARMACEUTICAL AMINES and quaternary ammonium compounds as ion pairs with anionic dye molecules has received a great deal of attention, investigations into the physical chemical basis of the observed phenomena have been relatively limited until recently. The extraction of quaternary ammonium compounds has been reviewed by Ballard et ai. ( I ) . Biles et al. have studied the partitioning behavior of amines paired with aromatic sulfonates and made an attempt to relate some physical parameters to the magnitude of the extraction constants (2, 3). Schill has recently reported on an extensive study of the extraction of amines and quaternary ammonium compounds with bromothymol blue (4-6) and also on the extraction of some high molecular weight amines with inorganic anions (7). A hypothesis based on the possible role of solvated ion pair species in enhancing the extractive process, proposed recently by Higuchi (2, S), assumes that the free energy involved in the (1) . , C . W. Ballard, J. Isaacs, and D. G. W. Scott. J. Pharm. Pharmacol., 6,971 (1954). ( 2 ) L. R. Hull and J. A. Biles. J . Pharm. Sci.. 53.869 (1964). . , (3j G. J. Divatia and J. A. Biies, Ibid.,50,916 (i961). (4) G. Schill, Acta Pharm. Suecica, 1, 101 (1964). (5) Zbid.,p. 169. (6) Ibid.,2, 13 (1965). (7) Ibid.,p. 99. (8) T. Higuchi and E. Roubal, Division of Analytical Chemistry, ACS, Abstracts of Papers, 149th Meeting, 28 B (April 1965).
transference of the ionic components from the water phase to form simple ion pairs in the organic phase would in many cases be far too unfavorable to yield useful partition coefficients. The addition of suitable solvating species to the organic phase would be expected to enhance markedly the extractive process through solvation of the formed ion pairs. Some evidence for this has already been presented by Biles (2). Experimental evidence presented here confirms the theoretical prediction that greater ion pair extraction results when suitable solvating agents are added to the nonpolar organic phase. It appears that formation of stoichiometric complexes between the ion pair and solvating species in the organic phase is more responsible for favorable distribution coefficients in most instances than are parameters such as dielectric constant (9). The studies were performed with dextromethorphan, d-3-methoxyN-methylmorphinan,
CH,O
as an example of a pharmaceutically important monoprotic organic base. GENERAL CONCEPT
The concept of the role of the solvating agent and its affinity for the ion pair has been considered by us in the following generalized manner. Ion pair solvation, or, equivalently, masking of the ionic character of the ion-ion bond, would be a significant factor in the extractive equilibrium if the ion pair structure was such as to expose strongly charged surfaces. These lipophilic ion pairs may, from a somewhat oversimplified viewpoint, be classified into three possible cases:
'
974
ANALYTICAL CHEMISTRY
case I
Case 11
Case 111
In the first case it is assumed that the cation is large and lipophilic except for the positively charged center. The small ex(9) P. Mukerjee, ANAL.CHEM., 28, 870 (1956).
posed anionic surface would be expected to carry a relatively high negative charge per unit area. This type of system may be effectively solvated by lipophilic molecules having an exposed positively charged surface-e.g., dipolar molecules with acidic protons such a s chloroform, phenols, and alcohols. Since the bonded solvating molecules would have their polar end buried adjacent to the anion, the appearance presented to the environmental solvent by the solvated ion pair is that of a relatively nonpolar aggregate. In the second case the situation is reversed, the ion pair having its cationic charge largely exposed. Solvating species containing nucleophilic sites may be expected to be particularly effective for this type of ion pair-e.g., ethers, ketones, amides, and phosphate esters. The third case is that of an ion pair with deeply buried charges. Having no exposed electrically unbalanced surface, it would be expected neither to require solvation to be readily extractable by nonpolar solvents nor be able to undergo appreciable solvation in theije media. The extraction of thr: ion pairs into the organic layer may be represented as done prwiously ( 2 ) by the equation: K.
BH+aq $. X-aq
$ BH+X-organic
(1)
where BH+,, represents the protonated base in the aqueous phase, X-sq the anion in the aqueous phase, BH+X-,,, the ion pair in the organic phsse, and K , the extraction constant for the equilibrium. A more rigorous expression must include the contribution made by the dipolar solvating agents. This may be expressed by the equation : BH+,,
+ X-*..,
+ n Mor.
[BH+X- * MJ,,,
(2)
-.Mor,
where M represents the solvating agent, and n, the effective molecularity of the solvating agent in the solvated ion pair. This expression will be governed by a n equilibrium constant, KO,which is defined as,: (3)
Experimentally the distribution ratio, D,is determined and can be represented as :
(4) Taking logarithms of Elquation 3, we have log KO = log D - log [X-]
- n log [MI
(5)
From this equation it is evident that at constant concentrations of the anion, a plot of log D us. log[M] should be linear, having a slope of n. From this plot the apparent stoichiometry of the association between the ion pair and the solvating agent may be readily evaluated. Since dextromethorphan essentially represents a cation of the type shown in Case I, where the positive charge appears to be shielded by a large lipophilic group and the anions studied appear relatively small in comparison, electrophilic solvating agents were investigated. Three solvating species representing three broad classes of proton donors were evaluated : halogenated hydrocarbons, the lipophilic alcohols, and the phenols, repres'ented by chloroform, 1-pentanol, and p-tert-butylphenol, respectively. In addition, a nucleophile, N,N-dimethylcaprylam de, was tested to confirm the need of a proton-donating agent.
MOLAR CONCENTRATION
OF Ne&
Figure 1. Effect of increasing bromide ion concentration in extraction of dextromethorphan pH = 3.0, 50x chloroform in cyclohexane as organic phase, p = 0.11, T = 25°C
EXPERIMENTAL
Equipment and Reagents. All absorbance readings were made on a Cary Model 11 M S spectrophotometer. Shaking was carried out on a Burrell Wrist-Action shaker. All temperatures were controlled to j ~ 0 . 2 ' C. Adjustments of pH were made on a Corning Model 12 Research pH meter. Dextromethorphan (obtained through the courtesy of Vick Divisions Research, Vick Chemical Co.) was found to be 99.7z pure based on nonaqueous titration. With the exception of NaBr, which was U.S.P. grade, and tropaolein, all chemicals were analytical reagent grade. Tropaolein was recrystallized from water four times. Water was prepared by distillation from acid permanganate in an all-glass distillation apparatus. The preservative was removed from chloroform by either distillation or shaking with concentrated sulfuric acid. The cyclohexane was spectrophotometrically pure. Procedure for Determination of Partition Coefficients. SINGLEEXTRACTION METHOD. Aqueous solutions were prepared in 250-1111 portions from stock solutions so that the final concentrations were 5 to 10 x 10-4M dextromethorphan, 0.001 to 0.10M NaX where X was the anion being studied, 0.01M H*O4-NaH2PO4 buffer, and sufficient &So4 to maintain a constant ionic strength. The pH was adjusted to 3.50 with HaP04. These solutions were prepared with water saturated with the organic solvents of the organic phase. The organic phases were prepared by mixing varying amounts of chloroform, n-amyl alcohol, and p-tert-butylphenol with cyclohexane and saturating the mixtures with water. Twenty milliliters of each phase were placed in a Teflonstoppered separatory funnel and shaken for 10 minutes. The phases were separated after standing for 5 minutes and the aqueous phase in each case was allowed to pass through a small plug of glass wool to remove minute globules of the organic phase suspended in it. The ultraviolet absorbance of each phase was then measured against a blank prepared in an analogous manner. From Beer's law plots prepared separately, the concentration of the dextromethorphan species in each phase was calculated and the distribution ratio was taken as
VOL. 39, NO. 8, JULY 1967
e
975
IC
a
2 .o
5.0
IO
[CHCI,] Figure 2. Effect of various anions on extraction of dextromethorphan as a function of chloroform concentration in cyclohexane pH = 3.00, p = 0.11, T = 25' C _ _ _ Iodide ........Ethylsulfonate -Bromide _ _ _ Chloride _ _ _ - Nitrate Concentrations. 5 X 10-4M dextromethorphan, 0.10M anion, 0.10M phosphate buffer. Line scalculated from experimental data to allow comparison
where E was the absorptivity of the dextromethorphan species in the respective phase. DOUBLE EXTRACTION TECHNIQUE.In cases where the KO's were found to be very high-for example, in the case of dextromethorphan picrate-it was necessary to employ a double extraction method to obtain reliable experimental results. In this technique an aliquot from the organic phase obtained from the first separation is equilibrated with fresh aqueous phase containing all constituents except the amine. Separation and spectrophotometric comparison of the amine concentrations in the two organic solvent fractions permitted convenient determinations of the partition coefficients in these systems. USE OF ACID-DYETECHNIQUE.When p-tert-butylphenol was used as a solvating agent, the spectrophotometric assay for dextromethorphan in the organic phase was not possible, because of interference. The acid-dye technique based on tropaolein 00 as described by Biles (3) was used for this reason to determine the concentration of the drug in both phases. RESULTS AND DISCUSSION
Chloroform as Solvating or Complexing Agent. Chloroform proved to be the simplest of the solvating species, since it does not undergo specific association in the organic phase. Chloroform-cyclohexane solvent systems were used for the bulk of the studies because they proved to be the least complicated of the organic solvent systems employed. 976
ANALYTICAL CHEMISTRY
LOG [CHClJ Figure 3. Effect of various anions on extraction of dextromethorphan as a function of chloroform concentration in cyclohexane pH = 3.00, p = 0.11, T = 25' C -0.10M trichloroacetate . . . . . . . . 1.OM benzenesulfonate 1 X lod3M picrate Concentrations = 5 X 10-4M dextromethorphan, anion concentrations as shown, 0.10M phosphate buffer, KSOa to adjust p. Lines calculated from experimental data to allow comparison
Figure 1 illustrates the typical change in the distribution ratio of dextromethorphan with the bromide ion concentration at constant pH, chloroform concentration, and ionic strength. In this case bromide ion concentration in the aqueous phase was varied, while the ionic strength was held constant by the addition of potassium sulfate. At pH 3.5 or less no significant amount of free base was extracted, as evidenced by the line in Figure 1 passing through the origin, only the ion pair being found in the organic phase. Sulfate anion was used to maintain a constant ionic strength, since this anion did not appear to form extractable ion pairs (IO). The influence of the nature of the anion on the apparent molecularity and distribution ratio, D,with chloroform as the solvating agent is evident in Figures 2 and 3. Data obtained for a number of organic and inorganic anions are shown plotted in accordance with Equation 5. The linear relationship between log D and log A4 seems to indicate adherence to the postulated relationship between the concentration of the solvating agent and the distribution of the ion pairs between the two phases. The slopes of the lines obviously correspond to the average number of molecules of chloroform apparently associated with each ion pair. In Table I experimentally determined values for n are listed together with calculated values for KO. Values for KOcannot be compared because the values for n vary from system to system. By calculating KOusing the equation,
(10) G. Schill, Acta Pharm. Suecica, 2, 119 (1965).
D 0.01
0006
I
I
I
0.40
0.60
I
0.80
1.0
[I - PENTANOL 1
0.5
Figure 5. Effect of 1-pentanol on extraction of dextromethorphan hydrobromide pH = 2.1, p = 0.11,T
I
If0
2.0
[I- PENTANOLI
Figure 4. Effect of 1-pentanol on extraction of dextromethorphan trichloroacetate pH
=
3.00, p
=
0.11, T
=
25' C
it is possible to list values that are independent of the value of n. These are also equivalent numerically to the value of KOat 1.O M solvating agent concentration. For anions having a highly exposed charged surface such as the halogens, nitrate, and the ethyl and benzene sulfonates, the number of chloroform niolecules bound per ion pair appears to lie in the range of 3 to 5 . In the case of the two organic anions, trichloroacetate and picrate, the data suggest that the resulting ion pair binds with a fewer number of the solvating molecules, probably because of the shielding effect of the organic portion of the anion. The picrate apparently becomes solvated by only two molecules of the proton donor, while the experimental value of n for the trichloroacetate was 2.36. Values of the equilibrium constant, KO,which can be calculated from the values of KO'given in Table I and Equation 6, appear to be highly dependent on the nature of the anion. This might be expected on the basis of their relative size and hydrophilic-lipophilic tendencies. This becomes apparent when the extractability cf identical anion ion pairs is compared in systems containing different solvating agents. 1-Pentanol as Solvatirig Agent. The solvating influence of an aliphatic primary alcohol on the extractability of an amine ion pair is evident in Figure 4. In this study an aqueous phase containing dextromethorphan and 0.1M trichloroacetate at pH 3.0 was equilibrated with an organic phase containing various concentrations of 1-pentanol in cyclohexane. The slope of t'ie straight-line log-log plot is approximately 1.50.
=
25" C
The apparent marked difference between the solvating behavior of chloroform and l-pentanol in similar systems can be rationalized on the basis that in organic systems the alcohol exists largely in dimeric and higher polymeric forms in these concentration ranges. If we assume that the average composition corresponds to that of the dimer, the binding number between the ion-pair and l-pentanol monomer would be roughly 3, which is comparable to the chloroform system. Results of a similar study on dextromethorphan-bromide ion system are summarized in Figure 5. The slope of the line shown corresponds to an exponential dependency of 1.96. This suggests that approximately four molecules of 1-pentanol were involved with each ion pair. This was less than that found with the chloroform system but still in general agreement with the previous findings. Solvation with ptertButylpheno1. The studies with the phenol which also exists in polymeric forms in the organic phase were treated somewhat differently. Because higher polymers of the phenol exist even at low concentrations, it was necessary to estimate the phenol monomer concentration in the organic phase in a different manner. To estimate
Table I. Calculated Extraction Constants for Ion Pair Extraction of Dextromethorphan and Various Anions and Apparent Binding Numbers Av. no. of CHC13 molecules associated with Anion KO,liters-mole-' each ion pair 4.9'8 Nitrate 4.0 x 10-4 4.94 6 . 0 x 10-4 Chloride 5.02 3.18 x 10-3 Bromide 4.20 Iodide 6.00 X 10-2 2.36 Trichloroacetate I 2.08 Picrate 5 . 6 X lo2 3.92 Ethyl sulfonate 3.6 X 4 . 1 X 10-2 4.82 Benzene sulfonate ~
~~~~
VOL. 39, NO. 8, JULY 1967
977
4 Figure 6. Distribution of p-tert-butylphenol between cyclohexane and water at 25" C
percent age conc. Lp-t -butyl phenol1 in the aqueous phase the amount of phenol monomer in the organic layer the distribution ratio for the phenol was determined for a series of phenol concentrations. This yielded the curve shown in Figure 6. When the curve was extrapolated to zero phenol concentration, the partition coefficient of the monomer was obtained and from this the monomer concentration for any total phenol concentration could be calculated. When dextromethorphan was partitioned in the presence of bromide at
varying concentrations ofp-tert-butylphenol, a plot of log D vs. log p-tert-butylphenol monomer yielded a straight line having a slope of 5.3 as shown in Figure 7. Again the dependence was similar to that found in the corresponding chloroformcyclohexane system. The effect of the change in solvating agent is summarized in Table I1 for several systems. The K l s have been calculated
n
n C
0.1
Ip-tertiary
I 04
t
I
.02
.03
I
os
butyl phenol monomer1
Figure 7. Effect of increasing p-rert-butylphenol concentration on extraction of dextromethorphan hydrobromide pH = 2.1, p = 0.21, T
978
ANALYTICAL CHEMISTRY
=
25" C
1
I
I
06
0.8
1.0
I
I
I
I
lolar Ratio Dimethycopr;%mid~~p-tertl'~ry butyl phenol. Figure 8. Influence of interaction of dimethylcaprylamide with p-tert-butylphenol on extraction of dextromethorphan hydrobromide pH = 2.1,
p =
0.26, T = 25" C
for 1M anion and 1M solvating agent concentrations. The varying magnitudes of these extraction constants may be attributed to the increasmg acidity of the proton of the solvating species in the order chloroform < n-amyl alcohol < p tert-butylphenol. Effect of N,N-Dimethylcaprylamide on Solvating Agent. It was of interest to determine if the schematic model postulated in Case I was reasonable. This was done by utilizing an electron donor solvating agent. This agent would not be expected to enhance partitioning of the ion pair. An experiment was carried out using fixed concentrations of dextromethorphan, bromide ion, and p-tert-butylphenol but adding increasing amaunts of dimethylcaprylamide to the organic phase. The results shown in Figure 8 indicate that the addition of the amide actually reduced the partition coefficient. These results would suggest that the amide did not act as a solvating agent even when present in rather high concentrations. Further, the decrease in partitioning at constant ptert-butylphenol could be expected, since in previous studies in these laboratories it l-as been shown that phenols formed strong association co nplexes with disubstituted aliphatic amides ( I I ) . Therefore, the addition of the amide resulted in the complexation of the phenolic solvating agent and essen(11) E. G. Shami, unpublished Ph.D. thesis, University of Wiscon-
sin. 1964.
Table 11. Effect of Various Solvating Agents on Extraction Constant for the Ion Pair Extraction of Dextromethorphan with Several Anions and Respective Molecularities Solvating Anion KO' Molecularity species 3 . 2 X 10-3 5.0 Chloroform Bromide 1.7 3.9 1-Pentanol 5 . 1 X lo7 5.3 p-tert-Butylphenol Trichloroacetate 7.0 2.5 Chloroform 51 3.0 1-Pentanol
tially prevented it from facilitating the ion-pair extraction. The requirement of a proton donor is further supported, since even when the amide was present in excess of the p-tert-butylphenol the nucleophilic agent did not enhance partitioning. ACKNOWLEDGMENT
We thank Smith Kline and French Laboratories, Philadelphia, Pa., and Warner Lambert Research Institute, Morris Plains, N. J., for support of these studies. RECEIVED for review July 11, 1966. Accepted April 10, 1967. Presented in part before the Symposium on Functional Group Analysis, Division of Analytical Chemistry, 149th Meeting, ACS, Detroit, Mich., April 1965.
Colorimetric Determination of Trace Levels of Oxygen in Gases with the Photochemically Generated Methyl Viologen Radical-Cation Philip B. Sweetser E . I . du Pont de Nemocrrs and Co., Wilmington, Del. A new colorimetric method has been developed for the determination of trace amounts of oxygen in gases by the in situ photochemical generation of reduced methyl viologen (l,l'-dimethyl-4,4'-bipyridiniumdichloride). The methyl viologen reduction is thought to take place by a photochemical electron transfer reaction initiated by the light excitation of a photoreceptor, proflavine (3,6-diaminoacridine), which induces the transfer of an electron to methyl viologen from an electron donor, EDTA.. The resulting deep-blue methyl viologen radical-catilon reacts rapidly with the oxygen to produce the original oxidized methyl viologen. The method is simple, rapid, and sensitive to less than 1 ppm oxygen. One of the outstanding features of this system is the elimination of cumbersome apparatus generally required for generating and maintaining the reducing agent.
IN MANY AREAS OF CHEMICAL RESEARCH, and in the chemical industry, there are numerous instances where it is important to measure the oxygen content of gases. Methods for oxygen analysis vary from a highly sensitive luminescent-bacteria1 method ( I ) , to colorinietric methods (2, 3), photonometric (1) K. P. Meyer, Helv. Phys. Acta, 15, 3 (1942). (2) L. W. Winkler, Ber., 21,2843 (1888). 20,1033 (1948). (3) L. J. Brody, ANAL.CF:EM.,
titrations of dissolved oxygen (4), dew point (5), galvanic cell-type methods (6, 7), a proposed oxygen meter based upon the photoreduction of methylene blue (8), and numerous others. Although many of the above methods are accurate and sensitive, most have some shortcomings which limit their general use. The requirement for maintaining a completely oxygen-free reagent necessitates the use of fairly cumbersome apparatus for generating and/or maintaining the reducing agent. The present paper describes a new colorimetric method which has been developed for the determination of trace amounts of oxygen in gases by the in situ photochemical generation of the reducing agent. The photochemical reducing system is based upon the coupling of Oster's (9, IO, 11) photoreducing dye system, proflavine and EDTA,
(4) T. Kuwana, ANAL.CHEM., 35, 1398 (1963). (5) L. Pepkowitz, Ibid., 27, 245 (1955). (6) P. Hersch, Nature, 169, 792 (1952). 25, 586 (1953). (7) M. G. Jacobson, ANAL.CHEM., (8) G. Oster and N. Wotherspoon, J . Chem. Phys., 22, 157 (1954). (9) G. Oster, Photographic Eng., 4, 173 (1953). (10) G. Oster and N. Wotherspoon, J . Am. Chem. SOC.,79, 4836 (1957). (11) F. Millich and G. Oster, Ibid., 81,1375 (1959). VOL. 39, NO. 8, JULY 1967
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