Chemical Stability
`of Pharmaceuticals
`A Handbook for Pharmacists
`
`Second Edition
`
`Kenneth A. Connors
`School of Pharmacy, The University of Wisconsin
`
`Gordon L. Amidon
`College of Pharmacy, The University of Michigan
`
`Valentino J. Stella
`School of Pharmacy, The University of Kansas
`
`A Wiley-Interscience Publication
`JOHN WILEY & SONS
`New York • Chichester • Brisbane • Toronto • Singapore
`
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`

`To the memory of Lloyd Kennon
`
`Copyright CD 1986 by John Wiley & Sons, Inc.
`
`All rights reserved. Published simultaneously in Canada.
`
`Reproduction or translation of any part of this work
`beyond that permitted by Section 107 or 108 of the
`1976 United States Copyright Act without the permission
`of the copyright owner is unlawful. Requests for
`permission or further information should be addressed to
`the Permissions Department, John Wiley & Sons, Inc.
`
`Library of Congress Cataloging in Publication Data:
`Connors, Kenneth A. (Kenneth Antonio), 1932-
`Chemical stability of pharmaceuticals.
`
`"A Wiley-Interscience publication."
`Includes bibliographies and index.
`I. Drug stability. I. Amidon, Gordon L. 11. Stella,
`Valentino J., 1946-
`. III. Title.
`[DNLM: 1. Drug
`Stability—handbooks. 2. Kinetics—handbooks.
`QV 735 C752cj
`
`RS424.C66 1986
`ISBN 0-471-87955-X
`
`615..18
`
`85-31455
`
`Printed in the United States of America
`
`10 9 8 7 6 5 4 3 2 I
`
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`

`CHAPTER 5
`Oxidation and Photolysis
`
`Oxidative and photochemical reactions are, for the
`most part, one-electron reactions as opposed to reac-
`tions discussed in Chapter. 4, which are two-electron
`reactions. For the hydrolytic reactions in Chapter 4,
`a free pair of electrons on a heteroatom in one mole-
`cule, a nucleophilic center, attacked an electrophilic
`center on a second molecule, whereas oxidative and
`photochemical reactions proceed through free radical
`or free-radical-like reaction pathways.
`Most drugs exist in a reduced form, so the presence
`of 20% oxygen in the atmosphere creates obvious poten-
`tial stability problems for these molecules. That is,
`many molecules tend to be converted to a more oxidized
`state. Kinetically, however, there is a sufficient
`energy barrier to many such reactions (the energy of
`activation) that not all molecules are subject to
`measurable rates of spontaneous oxidation or autox-
`idation. The radiation from the sun and artificial
`light, particularly visible and ultraviolet light, is
`also ubiquitous, so that molecules capable of rear-
`ranging upon absorption of radiation energy must be
`protected.
`Our overall mechanistic understanding of oxidative
`and photochemical reactions is poor. The reason for
`this will be understandable as this chapter proceeds.
`Simply stated, many oxidative and photochemical reac-
`tions involve very complex reaction pathways with mul-
`tiple intermediates so that even though the stoichi-
`ometry of a reaction might be given by Eq. (5.1) the
`kinetic law is not as simple as Eq. (5.2).
`
`RH + 0 2 ----> ROOH
`
`(5.1)
`
`82
`
`Oxidation
`
`83
`
`d[ROOHI
`dt
`
`k[RH][02]
`
`(5.2)
`
`Also, unlike two-electron reactions where catalysis is
`often limited to acid/base or nucleophilic catalysis,
`trace quantities of environmental agents can powerful-
`ly catalyze one-electron reactions. For example,
`trace contamination of metal ions can catalyze oxida-
`tive reactions by many orders of magnitude, and the
`presence of a photosensitizing agent can cause a mole-
`cule that in the absence of the photosensitizing agent
`is not photolabile to undergo an apparent photochemi-
`cal reaction.
`In this chapter we introduce, from a basic view-
`point, the kinetics and other factors affecting oxida-
`tive and photochemical reactions and describe how
`these reactions can be prevented or at least
`inhibited.
`
`A. OXIDATION
`
`1. Nature of Oxidation
`
`When one considers oxidation, it is important to real-
`ize that this reaction is a complementary one; its
`partner is reduction. One cannot happen without the
`other. Oxidation/reduction (redox) reactions involve
`the transfer of one or more oxygen or hydrogen atoms
`or the transfer of electrons. The classical, and
`familiar, inorganic redox system can be described by
`Eq. (5.3), where e- represents an electron and n the
`
`reduced form
`
`oxidized + ne-
`
`(5.3)
`
`number of electrons. Thus redox reactions are elec-
`tron-transfer processes, and this aspect must be con-
`sidered if the basic process is to be understood.
`In the case of organic compounds and especially the
`oxidation state of carbon, the oxidation state is de-
`termined by the number of bonds from carbon to oxygen.
`For example, the state of oxidation of one-carbon com-
`pounds increases as shown in Eq. (5.4). As stated
`earlier, the mechanism of this process is not as sim-
`ple as suggested by the stoichiometry of the reaction.
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`84
`
`Oxidation and Photolysis
`
`oxidation
`
`H
`
`H—C—H
`
`--> H -OH
`
`H.,
`-->
`C=0 --> H-C
`
`H
`
`H
`
`--* 0=C=0
`
`OH
`
`reduction
`Also the simple redox system illustrated by Eq. (5.3)
`is made more complex by the medium in which the reac-
`tion occurs. For example, the oxidation of hydroqui-
`none (1,4-dihydroxybenzene) to its quinone (p-benzo-
`quinone) is often illustrated in the textbooks by
`Eq. (5.5).
`
`(5.4)
`
`HO
`
`+ 2H+ + 2e-
`
`quinone
`hydroquinone
`Yet in aqueous solution, free electrons, e-, do not
`exist and the state of ionization of the hydroquinone
`is affected by the solution pH. Therefore in aqueous
`solution the oxidation of hydroquinones is more accu-
`rately described by Eq. (5.8).
`
`(5.5)
`
`al+
`
`+ H202
`
`+H20
`
`(5.6)
`
`As will be discussed later, the oxidation of hydro-
`quinone and other phenols is even more complex than
`shown by Eq. (5.6) in that the product of the immedi-
`ate oxidation, the quinone, can catalyze the oxidation
`
`Oxidation
`
`85
`
`further. This process is called autocatalysis or
`product. catalysis.
`
`2. Kinetics of Oxidation
`
`Oxidations that take place spontaneously under mild
`conditions are often called "autoxidation"; the major-
`ity of these are free-radical reactions. Free radi-
`cals are chemical species that possess an unpaired
`electron. Oxygen, in its ground state, is a diradical
`with the electronic configuration
`
`..6:•Cr•
`..
`
`Oxygen, therefore, would "like" to fill its outer
`electron shell to produce 022-, the peroxy dianion.
`To do so, oxygen must accept two electrons from a
`donor molecule(s) and in so doing could in theory
`generate other free-radical molecules.
`In most autoxidation reactions, even though oxygen
`is often involved, the initiation of the oxidation re-
`action does not involve molecular oxygen, that is,
`oxygen itself in its ground state does not really ini-
`tiate oxidation reactions. Let us consider the gener-
`al kinetic behavior of olefin autoxidation described
`by Eqs. (5.7)-(5.9)
`
`1(1
`R'' + R"-CH2-CH=CH-R"' --÷R"-all-CH=CH-R" + R'H (5.7)
`
`0-0'
`1
`k2
`R"-CH-CH=CH-R" + 02 -3 R"-CH-CH=CH-R"'
`
`(5.8)
`
`0-0'
`1
`k3
`R"-CH-CH=CH-R"' + R"-CH2-CH=CH-R"' --H>R"-6H-CH=CH-R"'
`
`(5.9)
`
`0-OH
`1
`R"-CH-CH=CH7R"'
`
`a hydroperoxide
`
`The reaction given by Eq. (5.7) represents the initia-
`tion reaction. Generally the species R'' is not
`
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`86
`
`Oxidation and Photolysis
`
`Oxidation
`
`87
`
`oxygen but some other peroxy radical present in the
`solution, trace quantities of metal ions such as
`ferrous or cupric ions, or radicals formed in the
`solution from the absorption of light (visible or
`ultraviolet).
`The reactions given by Eqs. (5.8) and (5.9) repre-
`sent chain propagation reactions, i.e., one radical
`produces one radical plus a hydroperoxide molecule.
`Radical species in solution, apart from reacting with
`oxygen or another unreacted molecule to produce anoth-
`er free radical, can also react with each other to
`produce stable or metastable products. This step is
`called a chain termination step and three examples are
`given in Eqs. (5.10)-(5.12).
`
`k4
`+ R02' --->
`
`R02
`
`'
`
`k5
`R02' + R' --->
`
`R' + R'
`
`k6
`--H›
`
`stable products
`
`Considering this reaction mechanism, and simplifying
`Eqs. (5.7)-(5.9) to Eqs. (5.13)-(5.15),
`
`initiator
`
`k 1
`
`R'
`
`R' + 02
`
`-. RO 2'
`
`R02' + RH
`
`k3
`
`ROOH + R'
`
`assuming normal levels of oxygen, applying a steady-
`state assumption to the radical species R02' and R',
`letting k 2[R'][02] = k3[R02'][RH], and k4 = k5 = k6 =
`kt, it can be shown that the rate of hydroperoxide
`formation is given by Eq. (5.16) where ri is the ini-
`tiation rate.
`
`d[ROOH] (ri kicni. Hi
`dt
`k t
`- -R
`
`k2[02]
`k3[RH] + k2[02]
`
`(5.16)
`
`It is obvious from this expression that the rate of
`hydroperoxide formation is proportional to the square
`root of the initiation rate, ri. Also, if oxygen
`
`concentration is very high, k2[02] >> k3[RH];
`Eq. (5.16) collapses to Eq. (5.17):
`
`d(ROOHj
`dt
`
`ri k
`) kq[RH]
`(kt
`
`(5.17)
`
`The reaction will apparently be first order in start-
`ing material, RH. On the other hand, if k3[RH] »
`k2[02], then Eq. (5.18) is realized.
`
`d[ROOHJ
`dt
`
`(ri )1/27,7 k2[02]
`
`(5.18)
`
`Under these conditions the reaction will be "pseudo
`zero order" with respect to RH, that is, [RH] does not
`appear in Eq. (5.18), and the reaction will be first
`order in [02]. If the term ri actually involves RH or
`02, the order of the reaction with respect to RH and
`02 could be as high as 1.5.
`The mechanism defined by Eqs. (5.7)-(5.15) assumes
`that the species ROOH is stable and that the termina-
`tion products are stable. In the case of olefins, the
`hydroperoxide can break down to produce volatile and
`nonvolatile products as well as multiple radicals.
`
`0-0H
`
`R"-CH-CH=CH-R"' --->R"-CHO + R"'-CH=CH' + 'OH (5.19)
`
`The rancidity of unsaturated cooking oils and oil-
`based paints is the result of this fragmentation of
`olefinic bonds to produce aldehydes, acids, and alco-
`hols as well as multiple radicals. As can be seen in
`Eq. (5.19), one hydroperoxide molecule produces two
`radicals. If this reaction is favorable, not only do
`we have a chain propagation reaction, but chain
`branching reactions will be observed. If branching
`does occur, the oxidation kinetics and products become
`even more complex than those defined by the initia-
`tion, propagation, and termination sequence.
`Qualitatively, the kinetics of oxidative free-radi-
`cal reactions follow the pattern illustrated in Figure
`5.1. A characteristic of many such reactions is a lag
`time or lag phase corresponding to the gradual buildup
`of radicals via the initiation step. If the radicals
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`

`S01 + MI"
`
`SOi. +
`
`SOi' + 02
`
`SOS.
`
`Oxidation
`
`89
`
`(5.20)
`
`(5.21)
`
`SOi •
`
`HS03-
`
`HS05 + SOi.
`
`(pH < 7)
`
`(5.22)
`
`-
`SO52 + 603 *
`
`(pH < 7)
`
`(5.23)
`
`HSO4 + SOg
`
`(pH < 7)
`
`(5.24)
`
`88
`
`Oxidation and Photolysis
`
`% Drug remaining
`
`-
`SOi. + SOS
`
`2 -
`SO3
`+ HSOi
`
`-
`-
`SO22
`+ SO2
`
`
`
`2- 2504
`
`(pH < 7)
`
`SO2. + SO;*
`
`, _2-
`02o6 + 02
`
`(5.25)
`
`(5.26)
`
`Time
`
`FIGURE 5.1. Illustration of percentage of drug re-
`maining vs. time for an oxidative free-radical reac-
`tion: curve a initiation step only; curve b initia-
`tion plus propagation; and curve c initiation, propa-
`gation, and chain branching. Arrows indicate the lag
`times.
`
`produced from the initiation step are stable then as
`soon as the catalytic species is consumed the reaction
`stops (curve a in Fig. 5.1). If the radicals produced
`from the initiator go into a propagation cycle, curve
`b results. The overall loss of drug will then often
`follow a first-order decay curve with respect to drug,
`depending on the oxygen dependency of the reaction.
`If chain branching occurs the overall loss of drug
`shows an acceleration phase (see curve c) with maximum
`acceleration occurring at -50% drug remaining.
`The reaction kinetics defined by Eqs. (5.7)-(5.19)
`were for a reaction in which the reactant RH was not
`capable of ionization. For the oxidation of drugs or
`pharmaceutical additives, the kinetics are further
`complicated when the state of ionization of the mole-
`cule is affected by solution pH. In the oxidation of
`sodium sulfite (or bisulfite) a first approximation of
`the oxidation mechanism is given by Eqs. (5.20)--
`(5.27),
`
`SOS' + inhibitor
`
`nonreactive products
`
`(5.27)
`
`where Eq. (5.20) is the initiation step, Eqs. (5.21)-
`(5.23) are propagation steps, Eqs. (5.24) and (5.25)
`oxidation steps leading to the ultimate oxidation
`product, S042-, and Eqs. (5.26) and (5.27) termination
`steps. 1.14- is a metal ion catalyst. The overall ki-
`netics of sulfate'formation or sulfite loss is very
`complicated, although it has been shown that at pHs <
`8 the proportionality given by Eq. (5.28) is observed
`(1).
`
`042-]
`a [M-41(s032-][11S03]
`d[Sdt
`
`(5.28)
`
`Note that in this expression there is no oxygen con-
`centration dependency, that the reaction proceeds
`faster at higher pHs, and that the reaction is sensi-
`tive to metal ion catalysis, especially by Fe2+, Mn 2+,
`and Cu 2+. The pH dependency arises because the frac-
`tions of sulfite and bisulfite ions are pH dependent.
`The kinetics and mechanism of oxidation of phenols
`and substituted o- and p-dihydroxybenzenes in aqueous
`solution are also very complex and pH dependent. In
`general terms the oxidation of such molecules is very
`sensitive to the presence of metal ions, oxygen con-
`centration, and pH, with increasing rates of oxidation
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`90
`
`Oxidation and Photolysis
`
`Oxidation
`
`91
`
`at higher pHs. This suggests that it is the anionic
`form of the phenol that is most sensitive to oxida-
`tion. In the case of dihydroxybenzenes, such as epi-
`nephrine and hydroquinones, the rate of oxidation
`often exhibits apparent second -order dependency on
`
`•
`
`OH
`
`+
`HO--p>"N— --NH2
`
`HO
`
`epinephrine
`
`hydroquinone
`
`The oxidation rates of
`hydroxide ion concentration
`hydroquinone and alkyl-substituted hydroquinones have
`been extensively studied because of their use in the
`photographic industry (2). Their mechanism of oxida-
`tion appears to be
`
`0
`
`semiquinone
`radical
`
`dimer
`
`H202
`
`0
`In both of the above reaction schemes it is the di-
`anion of the hydroquinone that appears to be the reac-
`tive species. The two pKa's of hydroquinone are >9
`and under pH conditions found in most formulations the
`fraction of hydroquinone present as its dianion is
`given by Eq. (5.29), where Kai and Ka2 are the first
`and second dissociation constants of hydroquinone.
`
`02
`
`+2H+
`
`H
`+ H20
`
`[HQ2- I
`f
`HQ2- - [HQ]
`TOTAL
`
`Ka 1Ka2
`
`[10]2 +
`
`+ Ka 1 Ka 2
`
`(5.29)
`
`HQ'H2
`
`H02
`
`However, it was found that the kinetics of oxidation
`of hydroquinone seemed to be dependent on quinone con-
`centration; that is, the quinone, the immediate prod-
`uct of the reaction, catalyzed the oxidation of the
`hydroquinone. This was explained by the following re-
`action, in which the quinone reacts with the dianion
`of the hydroquinone to form a very unstable semiqui-
`none radical, which very rapidly and spontaneously re-
`acts with oxygen to form two molecules of the quinone
`and hydrogen peroxide (H202), or dimerizes to a stable
`product (3)
`
`With the conditions [WE ] » Kai and
`becomes Eq. (5.30).
`
`Ka2
`
`Eq. (5.29)
`
`fHQ2-
`
`Kai Ka2
`
`Kai Ka2
`
` [OH- ]2
`
`[11+]2
`
`1(52„,
`
`(5.30)
`
`Therefore, if the rate of oxidation of the hydroqui-
`none is proportional to the hydroquinone dianion con-
`centration, it can be seen that the rate of oxidation
`will be apparent second-order in hydroxide ion concen-
`tration. This is confirmed by the data in Fig. 5.2
`for the oxidation of m-dimethylhydroquinone and hydro-
`quinone in aqueous buffer solutions. Similar pH de-
`pendencies on the rate of oxidation of ascorbic acid
`(see curves 3 and 4 of Fig. 1 of the L-ascorbic acid
`monograph) and captopril (see Fig. 1 of the captopril
`monograph) have been observed. In each case the rate
`of oxidative decomposition under aerobic conditions is
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`92
`
`Oxidation and Photolysis
`
`Oxidation
`
`93
`
`2.5
`
`- 3.0
`
`fp -3.5
`
`0
`
`- 4.0
`
`- 4.5
`
`70
`
`7.2
`
`7.4
`
`7.6
`pH
`
`7.8
`
`8.0
`
`8.2
`
`84
`
`pH -rate profiles for the oxidation of
`FIGURE 5 2
`hydroquinone (0) and m-dimethylhydroquinone (ED) at
`25°C. The slopes of the lines are 1.96 and 1.98, re-
`spectively. Both systems were studied in phosphate
`buffer.
`
`proportional to the fraction of the drug in the anion-
`ic form, namely the ascorbate anion or the thiolate
`(RS-) in the case of captopril.
`Interpretation of temperature effects on oxidative
`reactions is made difficult by the multiple steps in
`many of the reactions and because oxygen solubility in
`water (and other solvents) is temperature dependent.
`Since each reaction in a complex scheme will have its
`own activation energy, it is possible that as the tem-
`perature is changed a different reaction will become
`rate determining. Theoretically, under such circum-
`stances, the adherence of the reaction-rate/tempera-
`ture relationship to the Arrhenius equation will break
`down. Practically, however, over a limited tempera-
`ture range, Arrhenius behavior may be observed, but
`the activation energy is very much an "apparent" acti-
`vation energy for which the reaction conditions must
`be clearly stated. Included in this "apparent" acti-
`vation energy is the temperature dependence of the
`oxygen solubility. Table 5.1 gives the 02 content of
`water at various temperatures if the water is satu-
`rated by air or by pure oxygen. As can be seen for
`the air data, a 20°C change in temperature (5 —4, 25°C)
`results in a 40% decrease in oxygen concentration. If
`
`TABLE 5.1
`
`Oxygen Content of Water Under Air and Pure
`Oxygen at Atmospheric Pressure and Various
`Temperatures.
`
`Temperature (°C)
`
`Millimoles of 0 2
`from air per mL
`of 112011j 3
`
`Millimoles of 02
`from pure 02 per
`mL of H200,C
`
`0
`5
`10
`15
`20
`25
`50
`100
`
`0.386 x 10-3
`0.34 x 10-3
`0.304 x 10-3
`0.267 x 10-3
`0.232 x 10-3
`
`2.18 x 10-3
`
`1.29 x 10-3
`9.28 x 10-4
`7.51 x 10-4
`
`aFrom Reference 4.
`bCalculated from cc
`PV = nRT.
`cFrom Reference 5.
`
`of 02 in H2O and the expression
`
`the rate of reaction under study is first order in
`oxygen concentration, then a ninefold increase in rate
`(Q10 = 3) due to the direct effect of temperature on
`the rate-controlling step will show up experimentally
`as only a 5.5-fold increase in rate owing to the con-
`comitant change in oxygen concentration.
`
`3. Oxidative Pathways of Pharmaceutical Interest
`
`A few selected oxidative reactions of pharmaceutical
`interest are illustrated here; the stability mono-
`graphs include other examples. Many drug compounds
`have been reported to be subject to autoxidation, in-
`cluding adriamycin hydrochloride, amphotericin B, apo-
`morphine, ascorbic acid, cap'opril, chlorpromazine and
`other phenothiazine derivatives, cyanocobalamin, cys-
`teine, epinephrine, ergometrine, hydrocortisone, iso-
`amyl nitrite, isoproterenol, kanamycin, 6-mercaptopur-
`ine, morphine, neomycin, norepinephrine, novobiocin,
`p-aminobenzoic acid, paraldehyde, penicillin, phenyl-
`ephrine, physostigmine, prednisolone, prednisone,
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`94
`
`Oxidation and Photolysis
`
`procaine, resorcinol, riboflavin, streptomycin and
`dihydrostreptomycin, sulfadiazine, terpenes, the
`tetracyclines, thiamine, and vitamins A, D, and E.
`The unwanted conversion of fats, oils, flavors, and
`perfumes to a rancid state is due to oxidation of
`these unsaturated molecules. The double bonds are
`oxidized to form hydroperoxides, as demonstrated ear-
`lier, which then produce aldehydes; the latter cause
`the offensive odors and unpleasant flavors. Vitamin A
`(see monograph) and amphotericin B (see monograph) are
`two drug molecules with extended conjugated double
`bonds that are very susceptible to oxidative break-
`down.
`Epinephrine forms colored products on oxidation
`[Eq. (5.31)].
`
`HO
`
`H
`
`,CH3
`N,
`H
`
`HO
`
`epinephrine
`
`H
`
`,CH3
`
`02
`
`0
`
`epinephrine quinone
`
`melanin
`pigments
`
`-H20
`
`H3
`
`H3C H
`
`adrenochrome
`
`(5.31)
`
`The activity of riboflavin is contingent on its
`ability to take part in this redox equilibrium:
`
`H3C
`H3C
`
`CH2
`(CHOH) 3
`CH2OH
`
`riboflavin
`
`Oxidation
`
`95
`
`H3C
`
`H3C
`
`CH2 H
`(CHOW 3
`CH2OH
`
`dihydroriboflavin
`(5.32)
`
`Morphine dimerizes when oxidized. Equation (5.33)
`shows the first step.
`
`HO
`
`morphine
`
`HOO
`
`(5.33)
`
`The formed morphine free radical couples with a mor-
`phine molecule (at the free position ortho to the phe-
`nolic oxygen) to give the dimer (bimorphine or pseudo-
`morphine). Hydrogen peroxide is also produced and can
`cause additional oxidation to the N-oxide.
`Vitamin E can form an epoxide, which then produces
`a quinone: vitamin E (a-tocopherol):
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`96
`
`Oxidation and Photolysis
`
`Oxidation
`
`97
`
`monograph) is an example of a recently developed sulf-
`hydryl drug with oxidative stability problems. Sulf-
`hydryl-containing molecules generally oxidize to their
`corresponding disulfides, Eq. (5.36).
`
`R-SH
`
`Op R-S-S-R + H2O
`
`(5.36)
`
`4. Inhibition of Oxidation
`
`Exclusion of Oxygen
`
`If a molecule requires the presence of oxygen to de-
`grade why not exclude oxygen from the formulation?
`For parenteral drugs this can be achieved by packaging
`the drug in glass ampuls that are heat sealed under an
`inert atmosphere. For tablets, capsules, and so on,
`packaging of the formulation in ,a hermetic strip may
`be useful in preventing the oxidation.
`Quite often manufacturers would like to formulate a
`drug for parenteral administration in a multidose
`vial, which is sealed with a rubber stopper. It is
`very difficult to formulate very oxygen-sensitive
`drugs in such multidose vials because rubber is rea-
`sonably permeable to oxygen and both synthetic and
`natural rubber tend to release additives capable of
`catalyzing oxidative reactions. Oxidatively unstable
`drugs formulated in multidose vials require more than
`just oxygen exclusion during sealing to prevent oxida-
`tion from occurring.
`The importance of deoxygenating the water and the
`headspace atmosphere in an ampul prior to sealing can
`be seen in the following calculation using captopril
`as an example. The stoichiometry of the oxidative
`breakdown of captopril is given by Eq. (5.37), showing
`that 2 mol of captopril are lost for each half-mole of
`oxygen consumed.
`
`captopril disulfide
`
`(5.37)
`
`CH3
`
`vitamin E epoxide
`
`ck-tocopherylquinone
`
`(5.34)
`
`The phenothiazines readily oxidize, producing a
`multitude of products. The degradation mechanism of
`promethazine and the influence of pH, metals, chelat-
`ing agents, and antioxidants, have been extensively
`studied (6,7). The following reactions have been pro-
`posed and their products isolated.
`
`pH>4
`
`Sulfide and sulfhydryl-containing molecules are par-
`ticularly vulnerable to oxidation. Captopril (see
`
`(5.35)
`
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`98
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`Oxidation and Photolysis
`
`Oxidation
`
`99
`
`A 2 mg/2 mL injection of captopril, in a 2-mL ampul,
`has a headspace of -1 EL. The 2 mg of captopril is
`equivalent to 9.2 x 10-3 mmol of captopril. At 25°C
`there are 0.23 x 10-3 mmol of dissolved oxygen in each
`mL of water for a total of 0.46 x 10-3 mmol of oxygen.
`The 1 mL of headspace, assuming that it is air, con-
`tains 8.6 x 10-3 mmol of oxygen. Based on the stoi-
`chiometry of Eq. (5.37) it can be seen that the 9.06 x
`10-3 mmol of oxygen in the system would be capable of
`■ore
`degrading 3.62 x 10-2 mmol of captopril, that is,
`there is
`than enough oxygen present in this for-
`mulation to completely degrade the captopril. If the
`oxygen in the headspace is removed by flushing with an
`inert gas, the 0.46 x 10-3 mmol of oxygen in the water
`are still capable of degrading 1.84 x 10-3 mmol of
`captopril, or 20% of the formulation. By performing
`such calculations it is possible to predict how thor-
`ough the exclusion of oxygen from the system must be
`in order to prevent the oxidative breakdown of sensi-
`tive drugs.
`
`Alteration of Solution pH
`
`As has already been discussed, the oxidation of many
`drugs is pH sensitive. Acidic drugs such as ascorbic
`acid, phenols, and sulfhydryl compounds all degrade
`more rapidly in neutral to alkaline pH conditions.
`For such drugs the pH range 3 to 4 is generally found
`most useful in minimizing oxidation. Obviously, this
`pH range would not be useful for acidic drugs that
`have limited aqueous solubility at low pH values.
`Amine drugs such as the phenothiazines appear to be
`most stable in their protonated forms, that is, also
`at low pH values.
`
`Protection from Light
`
`Oxidative breakdown of drugs generally proceeds
`through the sequence of initiation, propagation (and
`maybe chain branching) , and termination. As mentioned
`earlier, a triggering force that may promote oxidation
`is "light" (namely, certain components of the electro-
`magnetic spectrum). Not all photolytic reactions are
`oxidative in nature and not all oxidative reactions
`require light as either an initiator or as an integral
`component of the propagation steps. If, however,
`
`light does initiate or promote an oxidative breakdown
`of a drug, the exclusion of the particular wavelength
`range of light responsible for the catalysis will
`often suppress the oxidation. This can be achieved by
`the total exclusion of light using an opaque container
`or the use of pigmented glass capable of excluding the
`damaging wavelengths. This will be discussed further
`in the section of this chapter on photolysis.
`
`Use of Chelating Agents and Antioxidants
`
`Oxidation reactions can be inhibited by agents that
`are
`
`chelating agents for metal ion initiators of
`free-radical oxidation reactions;
`reducing agents, that is, substances that can
`reduce an oxidized drug;
`preferentially oxidized compounds, that is,
`agents that are more readily oxidized than the
`agents they are to protect; or
`chain terminators, that is, agents capable of re-
`acting with radicals in solutions to produce a
`new species, a chain terminator radical, which
`does not reenter the radical propagation cycle.
`The new radical may be intrinsically stable or
`may dimerize to form an inert molecule.
`
`Compounds in all four categories are often classified
`as antioxidants (or antoxidants). It is probably more
`accurate to classify compounds in categories (b)-(d)
`as antioxidants and chelating agents as synergists.
`These categories are now treated in more detail.
`
`(a) Chelating Agents. Oxidative reactions are
`often initiated by metal ions such as Fe 3+ , Cu2+ ,
`Co3+ , Ni 2+, Mn 3+. These metal ions act as initiators
`because in their oxidation states they are capable of
`acting as radicals. For example, Cu2+ has 27 elec-
`trons and requires one more electron to complete the
`electron pair.
`Metal ions catalyze oxidative reactions in a number
`of ways. They can react directly with oxygen to pro-
`duce an oxygen radical, which can then initiate an
`autoxidation. The metal ion can form a complex with
`oxygen and subsequently form a peroxy radical. The
`
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`100
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`Oxidation and Photolysis
`
`Oxidation
`
`101
`
`metal ion can react with the drug itself to form a
`radical, as illustrated by Eq. (5.38), which is then
`able to enter into a propagation cycle [see Eqs. (5.7)
`and (5.13)].
`
`Mn+
`
`+ RH
`
`M(n-1)+ + 11 4- + ft'
`
`(5.38)
`
`The metal ion can also react with a hydroperoxide in
`the formulation to catalyze the breakdown, given by
`Eq. (5.39).
`
`M n+ + R'0011
`
`m(n-1)+
`
`H+
`
`R'02'
`
`(5.39)
`
`R'00H could be a hydroperoxide of the drug itself or
`of some other component of the formulation.
`Chelating agents act in an antioxidant capacity by
`binding metal ions, thus removing them, thermodynami-
`cally speaking, from solution. The most effective
`chelating agents used pharmaceutically are ethylenedi-
`aminetetraacetic acid (EDTA), citric acid, many of the
`amino acids, phosphoric acid (weak), and tartaric
`acid. EDTA and citric acid are the two most useful
`agents. Their metal-binding capacity is dependent on
`their state of ionization, both being most effective
`when their carboxylic acid groups are fully ionized.
`Thus, they lose their chelating capacity at low pH.
`
`- 00CCH2 \
`,CH2C00-
`/ N- CH2-CH2-N,
`-00CCH2
`CH2C00-
`
`p 2C00-
`
`HO-CH-000-
`
`CH2C00-
`
`EDTA tetraanion (edetate)
`
`citric acid trianion
`
`Just because an agent is able to chelate metal ions
`does not mean that it will reduce the effectiveness of
`the metal ion to act as a catalyst. There are circum-
`stances in which the metal ion may bind to some func-
`tional groups and in this bound capacity actually may
`be a better catalyst than in the unbound state. The
`chelating agents mentioned above, however, generally
`act to lower the catalytic activity of the metal ions
`towards radical chain reactions.
`
`(b) Reducing Agents. This approach is generally
`not used as a means of preventing oxidation. Sodium
`thiosmlfate and ascorbic acid are two reducing agents
`that have been used in this capacity.
`
`(c) Preferentially Oxidized Compounds. These are
`compounds that are more readily oxidized than the
`agents they are to protect. Essentially these agents
`act as oxygen scavengers. Two good pharmaceutical
`examples of oxygen scavengers are sodium bisulfite
`(and sulfite) and ascorbic acid. Sodium sulfite re-
`acts with oxygen according to Eq. (5.40).
`
`SO
`
`2-
`+ k0
`3
`2
`
`
`
`2-
`
`> SO
`
`4
`
`(5.40)
`
`It is possible to calculate the approximate amount of
`antioxidant needed to use up all the oxygen in an am-
`pul, for example, by calculating the amount of oxygen
`dissolved in the water and the headspace of the ampul.
`In this example, 2 mol of sodium sulfite equivalents
`(as sodium bisulfite) would be needed to consume each
`mole of oxygen.
`The sulfites are very commonly used, and a word of
`caution is in order. In the process of acting as an-
`tioxidants, sulfites yield acid sulfates, which cause
`a drop in pH. Also, sulfites can readily form inac-
`tive addition compounds, as with epinephrine (see
`monograph). They react with compounds such as al-
`kenes, alkyl halides, and aromatic nitro and carbonyl
`compounds. Sometimes, as with thiamine, they may
`cleave molecules.
`Although the nucleophilicity of the sulfite ions
`can be a disadvantage, there are circumstances where
`this nucleophilicity is an advantage. For example, in
`the oxidation of hydroquinone, it appears that the im-
`mediate product of the oxidation, its corresponding
`quinone, can act as a catalyst for further oxidation.
`Sodium sulfite (or bisulfite) can react with the qui-
`none and hydrogen peroxide to form hydroquinone mono-
`sulfonate, and sulfate as described by Eq. (5.41).
`
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`102
`
`Oxidation and Photolysis
`
`Oxidation
`
`103
`
`H20 2 + 2S03 2- ---*
`
`+ SO42- + OH
`
`OH
`
`(5.41)
`
`It appears that thiols can also act as antioxidants in
`a similar manner. All the major water-soluble antiox-
`idants, namely acetone sodium bisulfite, ascorbic
`acid, cysteine hydrochloride, isoascorbic acid, sodium
`bisulfite, sodium formaldehyde sulfoxylate, sodium
`metabisulfite, sodium sulfite, thioglycerol, thiogly-
`colic acid, and thiosorbitol, act as oxygen scaven-
`gers, although the thiol (sulfhydryl) group-containing
`antioxidants can also act as chain terminators (cate-
`gory d). Few if any of the lipid-soluble antioxidants
`act as true oxygen scavengers except for, perhaps,
`ascorbityl palmitate.
`
`(d) Chain Terminators. Referring to the earlier
`discussions on the oxidation of olefins [Eqs. (5.7)-
`(5.15)]. bisulfite [Eqs. (5.20)-(5.27)], and phenols,
`all of these reactions proceed through a radical mech-
`anism. Therefore any substance that can donate a hy-
`drogen radical while itself forming a radical that is
`stable and incapable of continuing the propagation
`chain cycle could act as an antioxidant. Such antiox-
`idants act by being acceptors. They are also called
`chain terminators.
`The major water-soluble antioxidants that can act
`as chain terminators are the thiol species cysteine,
`thioglycerol, thioglycolic acid, and thiosorbitol.
`These act by the mechanism described by Eq. (5.42).
`
`2R' (or R02' ) + 2R'SH --> 2RH (or ROOH) + R'-S-S-R'
`(5.42)
`
`Essentially all the lipid-soluble antioxidants act as
`chain terminators. The major examples are
`
`0
`
`HO
`
`hC15H31
`
`0
`ascorbityl
`palmitate
`
`HO
`
`6+
`
`OCH3
`
`butylated
`hydroxyanisole
`(BHA)
`
`CH3
`butylated
`hydraxytoluene
`(BHT)
`
`)._cH3T TH3
`
`H2CH-CH
`
`/>. -OH
`
`OH
`
`COOC3H7
`
`HO
`
`propyl gallate
`nordihydroguaia