Stability of Morphine in Aqueous Solution III
`
`Kinetics of Morphine Degradation in Aqueous Solution
`
`By SI-IU-YUAN YEH and JOHN L. LACH
`
`The degradation of morphine in aqueous solution is dependent on the pH of the
`solution and on the presence of atmospheric oxygen in the system. The overall
`reaction rate, in systems containing excess oxygen, was found to be equal to
`Kn
`H+
`.
`(
`>
`[at
`Kg + H+)] (Morphine)
`Degradation mechanisms for morphine, based on kinetic data and previously re-
`ported data on naphthol oxidation, are presented.
`
`Ka + H+
`
`MORPHXNE has been used as an analgesic and
`sedative since its isolation in 1805. Due
`
`to the limited solubility of morphine base, the
`acid salts, chiefly as the sulfate and hydrochlo-
`ride, have been used extensively in various
`pharmaceutical preparations.
`The stability of morphine in aqueous solution
`has been studied by many investigators since
`such solutions, after prolonged storage, undergo
`decomposition, as evidenced by discoloration.
`This decomposition of morphine is believed to be
`due to an oxidation reaction resulting in the for-
`mation of pseudomorphine (oxymorphine) and
`morphine N—oxide in the ratio of 9:1, together
`with a trace of a base believed to be methylamine
`(1). The oxidation of morphine and subsequent
`condensation to the dimer pseudomorphine
`is assumed to involve the phenolic group, as in
`the oxidation of naphthols to dimolecular com-
`
`pounds. Morphine derivatives not possessing
`the free phenolic group, as in the case of codeine
`and diacetyl morphine, do not undergo this type
`reaction (2). The oxidation of morphine is
`catalyzed by oxygen of air
`(3), sunlight
`(4),
`ultraviolet
`irradiation (5),
`iron and organic
`impurities (6), rat
`liver slices
`(7),
`tissue ho-
`mogenates (8), and cytochrome (9).
`Ionescu-
`Matin, at al. (10), claimed that the deterioration
`of morphine in presence of oxygen takes place
`through condensation of morphine at the phe-
`nolic hydroxy group with the formation of the
`dimer, pseudomorphine. However,
`in the ab-
`sence of oxygen, and in the presence of light,
`they also claimed that the deterioration is due to
`peroxidation or dimerization Which can take place
`through the oxygen of the hydroxy group which
`was thought to be activated by ultraviolet light,
`
`Received November 13, 1959, from the State University of
`Iowa, College of Pharmacy, Iowa. City.
`Abstracted in part from a dissertation presented to the
`Graduate School of the State University of Iowa by Shu-
`Yuan Yeh in partial fulfillment of the requirements for the
`degree of Doctor of Philosophy.
`
`35
`
`resulting in the formation of bimorphine. Abood
`and Kun (8) reported that in the course of oxi-
`dation of morphine by tissue, one mole of mor-
`phine utilizes one—half mole of oxygen and that
`the phenolic hydroxy group is oxidized to a
`quinone. Thorn and Agren (11) reported that
`the pseudomorphine formed in this oxidation
`was extremely stable and did not undergo further
`decomposition. However, Balls
`(12) pointed
`out
`that pseudomorphine was quite unstable,
`that it decomposes on either oxidation or re-
`duction, and that in alkaline solution pseudomor-
`phine gradually decomposed to higher oxidized
`products.
`The stability of morphine in aqueous solution
`is largely dependent on the hydrogen ion con-
`centration.
`In alkaline or neutral
`solution,
`
`morphine deteriorates rapidly at room tempera-
`ture, whereas acidic solutions are relatively
`stable (ll, 13, 14). The effect of temperature
`on the stability of morphine has been reported
`to be less important than the hydrogen ion con-
`centration (14).
`Although the stability of morphine in aqueous
`solution has been extensively investigated, no
`quantitative studies have been conducted. This
`report deals with a kinetic study of the degrada-
`tion of morphine in aqueous solution, the degrada-
`tion of these solutions carried out in a light proof
`oven. No study was made to investigate the
`efiect of light on morphine degradation.
`
`EXPERIMENTAL
`
`Reagents and Apparatus
`
`recrystallized from
`Morphine sulfate U. S. P.
`alcoholic aqueous solution and dried under vacuum
`for six hours, in. p. 250°; Beckman spectrophotom-
`eter, model DU, equipped with photomultiplier, 1-
`cm. silica cells. The buffers used in this study are:
`acetate bufier:
`0.2 M and 0.4 M at pH 3.0, 4.5,
`5.0, and 5.5; phosphate bufier:
`0.2 M and 0.4 M
`at pH 4.0, 5.0, 5.5, 6.0, 6.5, and 7.0. Phosphate
`
`   
`


`
`Lannett Holdings, Inc. LAN 1016
`
`

`
`QC
`
`Jommrl of Pharmaceutical Sciences
`
`_c
`
`I
`
`9
`~o\
`
`9
`x 85
`_.
`‘es-
`‘s
`‘3 N7
`'1‘,
`ti so
`.a
`
`N)“ __0
`
`0
`
`KO-
`
`‘\-Tc“-<3
`
`‘g
`E
`
`7.
`
`°
`
`3
`
`0
`
`°
`
`V
`
`\‘~
`
`1
`
`
`0.02 0.00
`
`_C_
`‘6L}
`
`75
`
`7U
`
`o
`
`i
`2
`
`absolute Miimgen
`Commercial Nilr’DqEIV
`
`4:.
`
`so
`
`40
`
`:20
`
`I50
`Time in *‘iLJurC
`
`200
`
`240
`
`280 300
`
`Fig. 2.AEffect of inner gas on degradation of mor-
`phine sulfate in deionized water, at 95°.
`
`04?
`
`0.36 -
`
`0.30
`
`0.24
`
`0.06 -
`
`004-
`
`Absorbonceat286my
`
`O
`
`20
`
`40
`
`60
`
`120
`100
`BO
`Time in Hours
`
`140
`
`160
`
`180
`
`Fig. 3.—Efi'ect of sodium bisulfite, sulfur dioxide
`on degradation of morphine sulfate at 95°:
`1 and 2
`contained 1% NaHSO;, sealed under atmosphere;
`3, M2804 in deionized H20, sealed under S02;
`1,
`assayed by direct spectrophotometric method;
`2
`and 3, by the chromatographic method.
`
`nitrogen showed no evidence of decomposition.
`The slight degradation noticed in systems saturated
`with commercial nitrogen was probably due to im-
`purities in the nitrogen, Absolute nitrogen was
`prepared by passing commercial nitrogen through
`three wash bottles containing Fieser’s solution (18)
`and through a saturated solution of lead acetate.
`It is interesting to point out here that morphine
`solutions underwent a color change to yellow when
`sulfur dioxide was introduced. This color intensity
`was dependent on the length of time used to saturate
`the solution with sulfur dioxide. Chromatographic
`analysis of
`these freshly prepared sulfur dioxide-
`saturated solutions gave low recoveries for morphine.
`Direct spectrophotometric analysis, however, of
`these same samples, using deionized water as the
`diluent, gave only slight absorbance difference when
`compared to the original solutions indicating that
`some reaction had probably taken place. These
`sulfur dioxide solutions become highly colored after
`three to four days’ storage at 95° and developed a
`yellow precipitate after one month. An investiga-
`tion of this phenomenon is, at present, continuing.
`It was also noted that
`the chromatographic
`
`36
`
`buffer, 0.2 M, at pH 2.0 and 2.5 were made by
`adjusting 0.1933 J11 phosphoric acid with mono-
`potassium phosphate. All buffer solutions were
`adjusted using a Beckman pH meter model H2.
`The phosphate buffer solutions used here are inde-
`pendent of the temperature (15), and the change of
`pH of acetate. bufier at higher
`temperature is
`negligible (16).
`
`Degradation of Morphine
`Sulfate in Sealed Ampuls
`
`Effect of pH on the Degradation of Morphine
`Sulfate Solution Sealed Under Atmosphere at 95°.!
`Five milliliters of 0.3% morphine sulfate solution in
`0.2 M phosphate buffer; pH 2.0, 6.0, 6.5. and 7.0,
`was introduced into 5-ml. ampuls, sealed under
`atmosphere, and stored in an oven at 95°. These
`solutions were assayed for morphine content at
`various time intervals by a previously described
`chromatographic procedure (17).
`The results of the study are shown in Fig. 1. The
`data obtained indicate that the rate and extent of
`decomposition of morphine is dependent on the pH
`of the solution.
`It is also apparent that after a
`time interval,
`this decomposition is halted, as
`evidenced by the plateaus. This behavior was
`probably due to the lack of oxygen in the systems.
`Calculation of the atmospheric oxygen present in
`the solutions contained in the ampuls (including the
`void space) was approximately 3.2 X 10“3 M/L.
`Since at pH 7.0 approximately 3.1 X l0“’ M/L. of
`morphine had undergone decomposition, it appeared
`that morphine and oxygen react on a mole to mole
`basis.
`
`*1
`
`95
`
`a:
`
`\J)
`
`.4 w
`
`5 5
`
`5 5
`
`
`
`McrpnineCm-.eim_iLi:r\M/L=4"J"
`
`
`
` %\ on 20
`
`o
`a*‘o—\__
`' **%3—~o——o——o_
`\t>\\‘_b&% PH 5 0
`
`.
`
`0
`
`an
`
`L
`an
`
`o\
`
`pH 6 5
`“o~——0——~5~ —o~——o}o
`pH 7 O
`— ——o
`
`4
`|6O
`:20
`lime in Hours
`
`1
`
`200
`
`i
`240
`
`_i
`290
`
`l.—EtTect of pH on degradation of mor-
`Fig.
`phine sulfate in 0.2 M phosphate bufi"er, sealed under
`atmosphere, at 95°.
`
`In order to verify the oxygen dependency of the
`reaction,
`the following study was undertaken.
`Five milliliters of 0.3% morphine sulfate solution
`was introduced into 5-ml. ampuls and sealed under
`sulfur dioxide, atmosphere, commercial nitrogen,
`and absolute nitrogen. An additional solution of
`morphine sulfate containing 1% of sodium bisulfite
`was prepared and sealed under atmosphere. Using
`sulfur dioxide and commercial nitrogen, the ampuls
`were bubbled for one minute,
`three minutes for
`absolute nitrogen, prior to sealing. These ampuls
`were stored in a 95° oven and assayed for morphine
`at periodic time intervals. Results of this study are
`shown in Figs. 2 and 3. The data indicated that
`the rate of degradation of morphine is oxygen»
`dependent since solutions sealed under absolute
`
`

`
`Vol. 50, No. 1, January 1961
`
`analysis of the morphine sulfate solution containing
`1% sodium bisulfite prior to scaling and storage at
`95° resulted in low recovery of morphine (70%).
`Direct spectrophotometric analysis of the freshly
`prepared solution, using deionized Water as the
`diluent gave, on the other hand, higher values for
`morphine. Further study dealing with this aspect
`showed there was no appreciable change in pH of the
`solution, although the peak of the spectrogram was
`shifted slightly to the lower wavelength from 286 to
`284 my and a stronger absorbance was noted, indi-
`cating that some reaction had taken place between
`morphine and sodium bisulfite. This phenomenon
`will be discussed in a future communication (19).
`This
`initial
`investigation confirmed previous
`studies and showed that
`the decomposition of
`morphine sulfate in aqueous solution was dependent
`on the hydrogen ion and oxygen concentration of the
`system. Attempts to employ a manometric pro-
`cedure using a Warburg manometer in the study of
`the oxygen dependency regarding the rate of degra-
`dation of morphine were unsatisfactory in that
`morphine sulfate, in 0.2 M phosphate buffer pH 7.0
`at 60°, underwent only slight color change after
`twelve hours, the amount of oxygen uptake being
`negligible. However, it was felt that Valuable in-
`formation with respect to the kinetics of this reac-
`tion could be obtained by maintaining a relatively
`constant oxygen concentration in the system. To
`this end, the rate of degradation of morphine was
`studied as a function of hydrogen ion concentration,
`molarity of buffers used, ionic strength, and concen-
`tration of morphine in oxygen-saturated systems.
`
`Degradation of Morphine Sulfate in the
`Presence of Excess Oxygen
`
`Effect of pH.—Approxirnately 0,15-Gm. portions
`of anhydrous morphine sulfate were accurately
`weighed, transferred to 100-1111. vaccine bottles, and
`dissolved in 50 ml. of the following buffers: 0.2 M
`phosphate bufier, pH 2.5, 6.0, 6.5, and 7.0; 0.2 M
`acetate bufier, pH 4.0, 4.5, 5.0, and 5.5. Three
`milliliters of each solution was Withdrawn and
`
`assayed chromatographically for the original mor-
`phine concentration. These solutions were bubbled
`with oxygen for two minutes, rubber stoppered,
`sealed with aluminum caps, and placed in a 95°
`oven. At various time intervals the bottles were
`removed, chilled, and 3-ml. aliquots of solution
`were withdrawn. After each sample removal, the
`remaining solution was again saturated with oxygen
`for two minutes, and stored in the oven. The proc-
`ess was repeated for a total of ninety-six hours.
`The data obtained, as shown in Fig. 4, indicated
`that
`the rate of degradation of morphine was
`hydrogen ion-dependent. A plot of the log of the
`concentration of undecomposed morphine against
`time gives a straight line, indicating that the reac-
`tion is pseudo first order with respect to morphine at
`constant hydrogen ion and oxygen concentration.
`A plot of the log of the specific rate constant of these
`reactions as a function of pH resulted in an “S”-
`shaped curve (Fig. 5). This plot resembles a dis-
`sociation curve and indicates that the rate of degra-
`dation of morphine is dependent on the type of
`morphine species present
`in the solution. Data
`obtained from studies conducted in 0.2 M phosphate
`buffer at pH 3.0, 4.0, 4.5, 5.0, 5.5, and 0.4 M phos-
`
`-2 O5
`
`37
`
`-215
`
`! v2.25
`
`’l
`
`LogMorphineConcentrationM/
`
`
`o
`20
`40
`so
`so
`100
`Time |l"I Hours
`
`Fig. 4.—-Efiiect of pHs on degradation of morphine
`sulfate in excess oxygen, at 95°.
`
`phate bufier at pH 5.0 were unsatisfactory in that
`the capacity of these buffers was inadequate.
`Effect of Buffer Molarity and Ionic Strength.—
`Data obtained in this study using 0.2 M and 0.4 M
`acetate buffer at pH 5.0 and 0.2 M acetate buffer at
`pH 5 containing 1 and 3% sodium sulfate indicate
`that the rate of degradation of morphine is inde-
`pendent of the molarity of buffer and of the ionic
`strength present, as shown in Figs. 6 and 7.
`Effect of Morphine Concentration.—-Although
`data already showed that the overall rate of degrada-
`tion of morphine at constant hydrogen ion and
`oxygen concentration is pseudo first order with
`respect to the concentration of morphine present,
`this first-order reaction was further verified by 3
`study of 0.3 and 0.15% morphine sulfate in 0.2 M
`acetate buffer, pH 5.0, as shown in Fig, 8,
`Effect of Temperature.~Solutions of morphine
`sulfate in 0.2 M phosphate buffer, pH 6.0, 6.5,
`0.2 M acetate buffer, pH 5.0, were subjected to
`degradation at 85, 90, and 95°. Results of this in-
`vestigation are shown in Figs. 9, 10, and 11. By
`plotting the log of specific rate constant, 12, against
`the reciprocal of temperature, T, as shown in Fig. 12,
`the apparent energy of activation, Ea, under the
`condition employed in this study, was calculated
`from the slopes of these lines and was found to be
`of the order of 22.8 Kcal.
`
`DISCUSSION
`
`Oxidation of Morphine.——A.s has already been
`pointed out, morphine undergoes decomposition
`resulting in the discoloration of the solution and the
`formation of precipitates. Our present study indi-
`cates that this degradation of morphine is chiefly
`dependent on the pH of the solution and the pres-
`
`

`
`38
`
`—l3
`
`—l5
`
`*1?
`
`L in
`
`I
`
`{V
`
`
`
`LogSpecificRateConstants .42 U1
`
`-2.‘)
`
`#2,?
`
`-29
`
`Fig. 5.—Plot of the log of specific rate constants as a
`function of pH.
`
`&\
`
`1
`
`02/‘.4
`0
`A 04/14
`
`‘ ~2,o5
`
`
`
`LooConrmnrouongrrv1o:pmng-M/L
`
`-1’\JmJa
`
`-2 EC
`
`0
`
`(C
`
`20
`
`30
`
`40
`
`60
`50
`TWTVE In Hours
`
`70
`
`80
`
`90
`
`I00
`
`Fig. 6.—-Efiect of molarity of acetate buffer at pH
`5 on degradation of morphine sulfate, in excess oxy—
`gen, at 95°.
`
`—?O6
`>203
`—2V0
`4:;
`.2 :4
`
`—2l6
`
`
`
`
`
`LoqMarnhheConcentrationM/L.
`
`r222
`
`o
`
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`
`20
`
`30
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`so
`so‘
`I-me ln Hqu"
`
`so
`
`70
`
`an’
`
`90
`
`:00
`
`Fig. 7.-—The effect of ionic strength on degrada-
`tion of morphine sulfate in 0.2 M acetate buffer, at
`pH 5.0 and in excess oxygen, at 95°.
`
`Journal of Pharmaceutical Sciences
`
`—2t5
`
`-2 E5
`
`Z3i
`
`m sW
`
`
`
`LogConcemrotnonanNlornhmzM/L
`
`015 “Q,
`
`o\‘ 0
`LogConcentrationofMorphineM/L,
`
`l0
`
`Z0
`
`50
`
`40
`
`60
`SD
`T.me m Hows
`
`70
`
`SO
`
`90
`
`V00
`
`Fig. 8.——Efi"ect of morphine concentration on deg-
`radation of morphine sulfate,
`in excess oxygen, at
`95°.
`
`-2.15
`
`-225
`
`-2.35 r
`
`-2.45
`
`0
`
`IO
`
`20
`
`30
`
`60
`50
`40
`Time IH Hours
`
`70
`
`80
`
`90 100
`
`Fig. 9.~Effect of temperature on degradation of
`morphine sulfate in 0.2 M phosphate buffer pH 6.0,
`in excess oxygen.
`
`the deterioration of morphine takes place
`that
`through a condensation process at
`the phenolic
`hydroxyl group with the formation of pseudomor-
`phine (10). Derivatives of morphine in which this
`phenolic group is alkylated, as in codeine, do not
`undergo this type of oxidation (2). The fact that
`deterioration of morphine increases rapidly in the
`presence of ultraviolet light (5) and decreases in a
`more acidic solution (11) gives additional support to
`the oxidation reaction. This oxidation of morphine
`to pseudomorphine may be somewhat analogous to
`the oxidation of phenols to dimolecular compounds
`and is presented here in some detail.
`The ferricyanide oxidation method commonly
`used to prepare pseudomorphine involves one elec-
`tron transfer Which is common in free radical sys~
`terns. This reaction can be illustrated as follows:
`
`It has been reported
`ence of atmospheric oxygen.
`(1) that the degradation products of morphine are
`pseudornorphine and morphine N-oxide together
`with traces of the base said to be methylamine, and
`
`[Fe(CN)5]"‘ + e -* [Fe(CN)5]“‘
`
`in their studies of the
`(20—23),
`Purnmerer, et al.
`oxidation of naphthols with alkaline potassium
`ferricyanide reported that the oxidizing agent attacks
`
`

`
`Vol. 50, No. 1, January 1961
`
`-2 05
`
`-—-
`
`-—
`
`4
`
`a
`
`39
`
`
`
`LeeContenlrotronofMorphmeM/L.
`
`-215
`
`-2.25
`
`-zssr
`
`i
`
`-245»
`
`T
`
`-2.55l-
`L
`
`-2 65.
`
`0
`
`0
`
`85°
`
`0900
`
`95°
`
`"
`
`‘
`
`1
`J
`
`1
`J
`
`.
`
`0
`
`IO
`
`20
`
`30
`F
`
`so
`so
`45
`Time in Hours
`
`70
`
`80
`J
`
`90 $00
`‘
`
`Fig. 10.——EiTect of temperature on degradation of
`morphine sulfate in 0.2 M phosphate buffer solution,
`pH 6.5, in excess oxygen.
`
`-2 05
`
`_
`371E‘
`Io X ii::“:T:::‘\“\—\\ so
`\.\\\
`.,.,\_\_;\\
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`-
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`-2 I5
`
`COHCENLFQIJD"oiMorphingM/L
`Log
`
`..
`{iv ‘
`
`
`
`—2 20
`
`30
`
`so
`40
`Time an Hours
`
`60
`
`‘lo
`
`so
`
`Fig. 11.—-Effect of temperature on degradation of
`morphine sulfate in 0.2 acetate bufier solution at pH
`5.0, in excess oxygen.
`
`the OH group and has a direct dehydrogenating
`action and that the primary oxidation product is a
`resonant aroxyl radical. They confirmed this by
`isolating and identifying the possible dimeric forms.
`The oxidation of naphthol gives binaphthol which is
`further oxidized to oxy-binaphthylene-oxide and
`binaphthylene dioxide as illustrated in Diagram I.
`Oxidation of p-cresol was found to give dicresol
`and an ether (24, 25).
`If the resonant forms of p-
`cresol are examined, it becomes evident that these
`are all cogent in the transition of the resonant aroxyl
`radical as illustrated in Diagram II.
`Considering such a reaction and by direct analogy,
`morphine could undergo a similar type of process in
`its degradation and subsequent production of pseu-
`domorphine as shown in Diagram III.
`Mechanism of Degradation of Morphine.——Since
`data obtained from this investigation showed that
`the rate of decomposition of morphine in solution
`was dependent on the presence of oxygen and that
`no decomposition occurred in systems void of oxy-
`
` LogSpecificRate
`
`Constants
`
`.04>
`
`.0on
`
`.0m
`
`0.7
`
`0.8
`
`o——
`2
`
`5*.
`3
`
`I
`4
`
`2.70 1
`
`0
`
`1
`5
`
`.3
`6
`
`/4jan__
`7
`8
`
`._
`
`92%|
`
`2
`
`% x 1,000
`12.——P1ot of the log specific rate constants
`Fig.
`as a function of the reciprocal of the absolute tem-
`perature.
`
`gen, it was concluded that a free radical reaction was
`involved in this process.
`It was also found that
`the rate of degradation was dependent on the
`hydrogen ion concentration of the solution since the
`rate was considerably greater at the higher pH.
`It
`is interesting to note here that an “S"—sha.ped curve,
`similar to the typical dissociation curve, was ob-
`tained when the log of the specific rate constant of
`these reactions was plotted as a function of pH.
`This indicated that the degradation was dependent
`on the type of morphine species present in solution.
`Based on this information, it appeared that the un-
`dissociated morphine molecules undergo oxidation
`more easily. At the pH values employed in this
`study and considering the pKa and pKb of morphine,
`only the protonated and undissociated morphine
`species are involved. The amount of anionic
`species of morphine present at these hydrogen ion
`concentrations is negligible and has been included in
`the concentration of
`the undissociated species.
`Since the rate of decomposition of morphine at
`lower pH’s, i. e., at pH 2.5 and pH 4.0, is nearly the
`same, this indicated that the protonated morphine
`species also undergoes oxidation, but at a diflerent
`rate to that of the free undissociated morphine base.
`The undissociated (free base form) and protonated
`morphine are both oxidized by atmospheric oxygen
`to give a semiquinone (M0) and a free radical per-
`oxide (HO;-). This semiquinone is further trans-
`formed to a free radical quinone (MO-), which can
`undergo coupling With:
`(a)
`itself,
`(b)
`the un-
`dissociated morphine, and (c) the protonated mor-
`phine. Since the amount of activated or free radical
`morphine species present in the system is small com-
`pared to the protonated or free base forms, inter-
`action or union of two such activated species is un-
`
`

`
`Journal of Pharmaceutical Sciences
`
`DIAGRAM I
`
`Oxidation of Naphthol (20-23)
`OH
`-
`
`O
`
`K,-,Fe(CN)5—w ~
`
`'
`
`O
`
`13
`
`OHOC OH
`
`————>
`
`GO OH2-»
`CO
`
`Oxy-binaphthylene Oxide
`
`Binaphthylene Dioxide
`
`DIAGRAM II
`
`Oxidation of p-Cresol (24, 25)
`CH3
`CH3
`
`0
`
`(pf?~Q«~¢
`
`O
`
`0
`
`0-
`
`O
`
`CH3
`
`Q?
`
`H
`
`40
`
`Interaction of this activated species with
`likely.
`the protonated or free base form of morphine is
`more probable resulting in the formation of the
`dimer pseudomorphine (PSM) with the simultaneous
`elimination of a hydr0gen—free radical H-. This
`hydrogen-free radical can then react with the per-
`oxide-free radical (HO2- ) to form hydrogen peroxide.
`The hydrogen peroxide which formed in such a
`process can react with morphine to form morphine
`N—0xide (MNO), or may decompose to give water
`and a free radical oxygen which can also react with
`morphine base to give the N-oxide. The postulated
`mechanism is given in Diagram IV.
`Derivation of
`the Rate Equation.~—~From the
`postulated mechanism and kinetic data obtained in
`this investigation the following rate equation may
`be derived
`
`M, = M + +HM = total concentration of morphine.
`M and +HM represent
`the undissociated and
`protonated form of morphine, respectively.
`MO and *‘HMO represent the activated forms of
`the undissociated and protonated morphine
`species, respectively.
`
`k
`M+O2—:MO- +HO2-
`k
`+HM + 02 -3 +HMo- + H02’
`12
`+HMO- + M—3>PSM + H-
`1;
`+HMo- + +HM—4>PSM + H-
`Ie
`MO-+M—:PSM-I-H»
`k
`MD» + +HM—G>PSM + H-
`
`161
`H‘ + H02’ '—* H202
`k
`H202 + M -5 MNO + H20
`
`Application of the usual steady-state treatment
`for elimination of the unstable free radical and inter-
`mediates gives the rate.
`
`—d(Mc)/dz = 3k1(O2)(M) + 3k2(O2)(+HM)
`
`The dissociation equation of an ampholyte may
`be represented as follows
`
`"'HM—->M+H"'
`
`Ka
`
`= <M><H+>
`<+HM>
`
`M,=M++HM
`
`Ka
`M “ M‘ Ka+H+)
`
`.HM = M‘ Ka + H’'
`Therefore, the overall rate can be written as
`
`CH3
`
`CH3
`
`+
`
`———,
`
`OH
`
`0'
`
`CH3
`‘CH3
`O O ’\ l 0
`
`OH
`
`OH
`
`OH
`
`Since oxygen was maintained in excess, this rate
`equation becomes pseudo first Order.
`
`~d(Mz)/di =
`i’**’ (ifiaiel + t’ (mififill ‘M9
`When a Ka Value of 1.7 X 10-7 was chosen for
`morphine, the Values of k1’ and kg’ were calculated
`and found to be 8.03 X 10"’ per hour, and 1.77 X
`10‘3 per hour. The success of this derived equation
`is shown in Table I.
`It is interesting to note the
`close agreement of the Values obtained using this
`overall rate equation as compared to the reported
`experimental data.
`
`CONCLUSION
`
`-d(M:)/dz‘, = [sk,(o,)
`
`+ 3/e2(O2)
`
`The degradation of morphine in aqueous solu-
`tion is dependent on the pH of the solution and
`on the presence of atmospheric oxygen in the
`
`(Mg
`
`

`
`Vol. 50, No. 1, January 1961
`
`41
`
`DIAGRAM III
`
`Formation of Pseudomorphine
`
`N*—CI-I3
`
`DIAGRAM IV
`
`Degradation of Morphine
`N-CH3
`
`
`
`OH
`
`-O
`
`OH
`
`N-CH3
`
`O
`
`OH
`
`N—CH3
`
`N_CH3
`
`' 0
`
`O
`
`OH
`
`Q
`0 U +
`
`O
`
`.
`
`O
`
`->
`
`OH
`
`OH
`
`O
`
`OH
`
`N—CHa
`
`N—CH3
`
`N*—CH3
`«
`O ‘ +
`
`O
`
`OH
`
`.
`
`O
`
`O
`
`N—CH3
`“’
`
`+ H_
`
`0
`
`OH
`
`OH
`
`O
`
`OH
`
`OH
`
`H‘ + H02 H’ H202
`
` N—CH3
`
`O
`
`O
`
`+ H202 —__.+
`
`OH
`
`OH
`
`O TN
`
`'"CH3
`
`0
`
`O
`
`+ H202
`
`OH
`
`OH
`
`TABLE I.-“OBSERVED AND CALCULATED SPECIFIC
`RATE CONSTANTS AT VARIOUS pH’s
`
`21)};
`4:0
`4.5
`5.0
`2'2
`7:0
`
`5.5
`
`I: X 103 l11'.'1,
`Ohierggd
`1:95
`2.16
`2.99
`$33’?
`48‘-70
`
`5.61
`
`la x 103 hrfl,
`Calcgl. by the
`Derive./97Eq'
`1:90
`2.19
`3.08
`536%
`5120
`Mean aw"
`
`5.77
`
`Deviation,
`+87"5
`_2‘_5
`+1.4
`+3.0
`:3‘?
`+5-1
`0'62
`
`+2.8
`
`in systems
`system. The overall reaction rate,
`containing excess Oxygen, was
`found to be
`e 113] to
`C1
`K
`H+
`I“ (1%) + k2 (K+T+>] ‘M°’Ph‘“”
`The apparent energy of activation, Ea, was
`calculated to be 22.8 Kcal. for this reaction.
`
`a
`
`,
`
`.
`
`

`
`42
`
`Journal of Pharmaceutical Sciences
`
`REFERENCES
`
`(1) Kollo, C., Bull. soc. chim. Romania, 1, 390919).
`(2) Vongerichten, E., A7m,, 294, 2060896).
`(3) Lesure, A., J. pizarm. clu'm., (6) 30, 337(1909).
`(4) Kiss, M. M., Ber. zmgar. pharm. Ges., 16, l28(1940).
`(5) Buchi,
`_l., and Welti, H., Pharm. Acta .Hel'v., 16,
`e7(1941).
`_
`(6) Balan, J., and Csere, E., Chem. Wash. 1, 407(1953).
`(7) Betnheim, H., and Bernheitn, M. L. C., J. Pharmacol.
`Exptl. Therajx, 81, 374(l944).
`(8) Abood, L. G., and Kun, E_, Federation Proc., 18,
`207(1949l.
`(9) I-Iosoya, E., and Brody, T. M., J. Pharmacol. Exptl.
`rlmap, 120, 505r1957).
`(10)
`Ionescu-Matiu, A., Papescu, A., and Moncium, L.,
`Ann, Pharm.f1anc., 6, 137(1948).
`(11) Thom, N., and Agren, A.. Swans}: Farm. Tidskn, 55,
`61(1951).
`(12) Balls, A. K., .7. Biol. Chem., 71, 537(1927).
`(13) Zoccola, 17., Giant, farm. chz'm_, 67, 60, 900918).
`(14) Dietzel, R., and Hues, W., Arch. Pharm., 266,
`641(1929) .
`
`(15) Bower, V., and Bates, R. C., J. Research Natl. Bur.
`Standards, 59, 261(1957).
`(16) Kolthofl, I, M., "Acid—Base Indicators," The Mac-
`millan Co., New York, N. Y., p. 265.
`(17) Yeh, S. Y,, and Lach, J. L., to be published.
`(18) Fieser, L. F., “Experiments in Organic Chemistry,”
`3rd rev. ed., D. C. Heath & Co., Boston, Mass, 1955, p, 299.
`(19) Yeh, S. Y., and Loch, J. L., to be published.
`(19
`.
`Putumerer, R., and Frankfurter, F., Ben, 47, 1472
`(21) Pummerer, R., and Cherbuliez, E,,
`ibid., 47, 2957
`(1914).
`(22) Pummerer, R., and Cherbuliez, E.,
`1402, 1403, 1414(1919).
`(23) Pummerer, R., and Frankfurter, F,,
`(1919).
`(24) Pummerer, R., Melamed, D., and Puttfarcken, H.,
`ibid-'., 55, 3116(1922).
`(25) Pummerer, R., Puttfarcken, 1-1., and Schopflecher, P.,
`ibid.',, 58, l808(1925).
`
`ibid., 52, 1416
`
`ibid., 52, 1392.
`
`Antifungal Activity of Some Amides of
`Dichloroacetaldehyde and Bromal
`
`By WILLIAM D. EASTERLY, jr., and JAMES E. DUSENBERRY
`
`The fungus-inhibiting properties of a series of amides of dichloroacetaldehyde and
`bromal were studied. Three of the compounds showed antifungal activity when
`subjected to tests with Aspergillus niger and Trio/Joderma viride. Terminal chlorine
`atoms appeared to be a factor in enhancing the antifungal activity.
`
`RECENTLY the syntheses of some amides of
`dichloroacetaldehyde ( 1) and bromal
`(2)
`were reported.
`In View of the fact that studies
`have shown the potentialities of certain sub-
`stituted amides as antifungal agents and have
`indicated an increase in antifungal activity in
`certain compounds following halogenation (3, 4),
`it was decided to subject
`these compounds to
`
`preliminary tests.
`
`EXPERIMENTAL
`
`the fungusinhibiting
`Procedure.—A study of
`properties of these compounds was carried out by a
`modification of a method used by Bateman (5),
`Vincent (6), and also by Leonard and Blackford (3).
`This procedure consists in comparing the growth
`rates of the test fungus upon nutrient agar contain-
`ing known concentrations of the compound to be
`tested with that of a control, identically treated but
`containing none of the test compound. All com-
`pounds were tested in triplicate for each concentra-
`tion.
`Preparation of Culture Medium.—The culture
`medium, Sabouraud’s dextrose agar, was prepared
`
`Received July 18, 1960, from the University of Arkansas,
`School of Pharmacy, Little Rock.
`Accepted for publication August 4, 1960.
`Supported by grants from the Upjohn Co. and the Sigma Xi
`Society.
`
`in the usual manner; and prior to being steam
`sterilized for twenty minutes at 15 pounds pressure,
`9. glass enclosed magnetic stirring bar was added to
`the flask containing it. When the flask was removed
`from the autoclave it was placed in a hemispherical
`heating mantle, which in turn was placed on a mag-
`netic stirring motor. A weighed amount of the test
`compound was added to the agitated culture medium
`which was maintained in the liquid state.
`Most of the amides of dichloroacetaldehyde, in the
`indicated concentrations, were readily soluble in the
`culture medium, while the amides of bromal were
`difiicult to dissolve. The latter could be uniformly
`dispersed, however, by keeping the magnetic stir-
`ring bar moving rapidly, and gradually lowering the
`temperature of the heating mantle until the con
`sistency of the culture medium was still suitabie for
`pouring, yet keeping the test compound in suspen-
`sion. The culture medium was then poured into
`previously chilled, sterile 9-cm. Petri dishes.
`Inoculation.——The test organisms to be used had
`been grown on Sabouraud’s agar slants for seven
`days at 30°
`Separate spore suspensions of each
`of the two test organisms were prepared by washing
`the agar slants with a 5-ml. portion of normal saline.
`As an inoculum, 0.03 ml. of this spore suspension
`was added to the center of the Petri dishes contain—
`ing test substance, and a. like amount was added to
`the control.
`Growth Measurement:-—The plates, incubated at
`37°, were observed daily and measurements of