factors may affect the overall dissolution profiles such as deviation from
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`the log-normal law, irregularity in shape, and differences in the diffusion
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`barrier for each particle, none of which is easily available. But these
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`factors do not prevent an understanding of particle-shape effects on drug
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`dissolution profiles, since they normally act independently of the ef-
`facts.
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`In practical terms, the sizes of acicular or flaky particles measured
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`
`microscopically or by an automated counter tend to be larger than those
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`available for the evaluation of their dissolution profiles since the micro-
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`scopic method does not always give the smallest side length but gives a
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`larger one and the automated counter method gives volume diameter.
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`REFERENCES
`
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`(1) W. I. Higuchi and E. N. Hiestand, J. Pharm. Sci, 52, 67 (1963).
`
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`(2) W. I. Higuchi, E. L. Rowe, and E. N. Hiestand, ibid., 52, 162
`(1963).
`
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`
`
`(3) D. Brooke, ibld., 62, 795 (1973).
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`(4) Ibid., 63, 344 (1974).
`
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`(5) P. V. Pedersen and K. F. Brown, J. Pharm. Sci., 64, 1192
`(1975).
`
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`(6) J. T. Carstensen and M. Patel, ibid., 64, 1770 (1975).
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`('7) P. V. Pedersen and K. F. Brown, ibid., 65, 1437 (I976).
`
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`(8) “Handbook of Mathematical Functions with Formulas, Graphs,
`
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`and Mathematical Tables,” M. Abramowitz and I. A. Stegun, Eds., Na-
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`tional Bureau of Standards, Washington, D.C., 1965.
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`otki
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`ski
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`Figure 2——Rectangular parallelepiped having the dissolution rate
`
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`constants ki, aki, and /ikg.
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`average dissolution rate constant, h,-,. (M = 1.0, 0' = 0.5, p = 1.5, a = 1.0,
`
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`
`
`and )3 = 10.0), and with isotropic behavior are presented in Table Ill. Both
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`profiles are almost the same. The ratio of the time necessary for 50%
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`dissolution is 120.95. This result implies that the evaluation of the dis-
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`solution for the acicular or flaky particles with nonisotropic behavior is
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`roughly possible by means of Eq. 6 using the mean dissolution rate con-
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`stant, km, which may be obtained experimentally as described.
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`The basic assumptions behind the theory are that the constituent
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`particles are rectangular parallelepipeds that are similar in shape and
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`dissolve isotropically under sink conditions. In actual situations, more
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`Autoxidation of Polysorbates
`
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`M. DON BROW X, E. AZAZ, and A. PILLERSDORF
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`Received September 6, 1977, from the Pharmacy Department, School of Pharmacy, Hebrew University, Jerusalem, Israel.
`
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`publication March 31, 1978.
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`
`Accepted for
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`Abstract 0 Aqueous solutions of polysorbate 20 undergo autoxidation
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`on storage, with the peroxide number increasing and subsequently de-
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`creasing again, the acidity increasing continuously, the pH and surface
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`tension falling and tending to level off, and the cloud point dropping
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`sharply until turbidity begins at room temperature. The changes are
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`accelerated by light, elevation of temperature, and a copper sulfate cat-
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`alyst. At the same time, hydrolysis occurs, liberating lauric acid. Analysis
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`of the alterations in these properties leads to the conclusion that hy-
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`drolysis has the major influence near room temperature and that oxy-
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`ethylene undergoes chain shortening at temperatures above 40°. How-
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`ever, evidence of degradation is detectable even in previously unopened
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`commercial samples of polysorbates 20, 40, and 60, warranting attention
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`to the stability of and standards for these surfactants as compared with
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`the solid alkyl ether type of nonionic surfactant.
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`Keyphrases D Polysorbates, various—autoxidation on storage, effect
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`of light. temperature, and copper sulfate U Oxidation—various poly-
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`sorbates on storage, effect of light, temperature, and copper sulfate El
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`Stability—various polysorbates. autoxidation on storage, effect of light,
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`temperature, and copper sulfate [:1 Degradation—various polysorbates,
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`autoxidation on storage, effect of light, temperature, and copper sulfate
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`:1 Surfactants—various polysorbates, autoxidation on storage, effect of
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`light, temperature, and copper sulfate
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`In view of accumulating evidence of the ease of autox-
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`idation of polyethylene glycols and polyoxyethylene fatty
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`alcohol ethers (1-4), it was suspected that other nonionic
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`surfactants might undergo a similar process. Information
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`about such reactions could increase the understanding of
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`drug instability in aqueous solutions containing nonionic
`
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`surfactants (1, 3).
`
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`
`The only systematic investigation of autoxidation in
`
`
`
`1616 / Journal of Pharmaceutical Sciences
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`Vol. 67, No. 12, December 1978
`
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`nonionic surfactants was carried out on cetomacrogol (1,
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`2). Peroxides were formed and decomposed spontaneously
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`at rates increasing with temperature and decreasing with
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`surfactant concentration. Furthermore,
`the induction
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`period for peroxide chain propagation was shortened by
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`an increase of temperature, a reduction of pH, a copper
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`sulfate catalyst. The period was also reduced by the ad-
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`dition of chemical initiators, such as hydrogen peroxide
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`or partially oxidized surfactant, and by free radical-ini-
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`tiating processes, such as exposure to light or thermal
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`treatment as in sterilization by autoclaving.
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`During storage, the pH and cloud point fell and the acid
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`content rose while the surface tension characteristics
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`changed drastically. Polyglycols exhibited parallel changes
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`in peroxide and acid content and in pH after autoclaving.
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`These changes were interpreted as showing that degra-
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`dation occurred in the hydrophilic chain with progressive
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`reduction of the oxyethylene content until the hydro-
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`phi1ic—lipophilic balance fell below the critical value for
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`solubility in water, when phase separation of the surfactant
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`occurred at room temperature.
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`Since sorbitan derivatives are used widely, knowledge_
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`of their stability is important. Their behavior relative to
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`the fatty alcohol ether type of surfactant may govern the
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`choice between these two agents in a formulation. In the
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`present work, the decomposition of polysorbate 20 (poly-
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`oxyethylene 20 sorbitan monolaurate) was studied sys-
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`tematically at controlled temperatures. The results are also
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`ao22-3549/ 73/ 1200-1676$01.00/0
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`© 1978. American Pharmaceutical Association
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`SENJU EXHIBIT 2097
`
`Page 1 of 6
`
`SENJU EXHIBIT 2097
`LUPIN v. SENJU
`IPR2015-01105
`
`

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`DAYS
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`Figure l—-Rate of peroxide formation in 3% aqueous polysorbate 20
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`at 70, 60,40, and 25° in daylight (a), darkness (b), or daylight with I X
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`10"‘ M copper sulfate catalyst (c). P.N. = peroxide number in millie-
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`quivalents per kilogram.
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`rate is or has been high. Therefore, a low P.N. value is not in itselfa cri-
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`terion of nondegradation of the surfactant.
`
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`Surfactant Chain Degradation—Various chemical degradation
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`products have been detected, including carbonyl compounds and acidic
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`products (12), but no stoichiometric relationship with the degree of de-
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`composition of the surfactant has been established. Acid formation is
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`readily measured by the change in pH and the total acid content, and it
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`is an important indicator of the extent of degradation and the quality of
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`the surfactant (1, 12). Increased degradation leads to larger amounts of
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`low molecular weight acids and, thus, a lower pH. The terminal hydroxyl
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`groups are relatively stable compared with hydroperoxide and free rad-
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`icals, requiring highly acidic conditions with strong oxidizing agents for
`
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`their oxidation to terminal aldehydic and carboxylic groups.
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`However, substitution of a hydroperoxide radical in the a— or (3-position
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`to the terminal hydroxyl leads to instability, which generally causes mi-
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`gration and/or C—C or C-O fission (9, 11, 12) with consequent formation
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`of a two-carbon acid or formic acid, respectively. The latter has been
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`identified as a degradation product (12).
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`In view of the complexity of the degradation reactions and the difficulty
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`of performing separations quantitatively in dilute aqueous surfactant
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`solutions, physical-chemical methods have particular importance. Sur-
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`face tension—concentration curves enable measurement of changes in the
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`critical micelle concentrations (CMC); furthermore, the sub-CMC slope
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`is determined by the surfactant area per molecule at the air—liquid in-
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`terface if the molecules are close packed. In the oxyethylene type of
`
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`nonionic surfactant, the area per molecule at close packing is determined
`
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`by the number of oxyethylene groups present in the molecule and is rel-
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`atively independent of the hydrophobic group due to coiling of the
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`polyglycol groups (13). The surface tension value above the CMC is rel-
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`atively constant and is also a characteristic property, becoming pro-
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`gressively lower in a series based on the same hydrophobic group as the
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`hydrophilic chain is shortened (2).
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`The cloud point of nonionic surfactants is the temperature at which
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`turbidity appears on heating their aqueous solutions. It is related to the
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`hydrophilic-lipophilic balance; in a series based on the same hydrophobic
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`group, phase separation occurs at progressively lower temperature as the
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`oxyethylene chain length is reduced (2), but the cloud point is also sen-
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`sitive to the presence of additives (14).
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`Changes in surface tension and cloud point properties were used to
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`establish that degradation occurred in the hydrophilic chain in ceto1na~
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`crogol and to estimate the rate of loss of oxyethylene groups at 530°
`(2).
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`For polysorbate 20, chain breakdown may be represented as shown in
`Scheme Vl:
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`C11H23c0(OR) (OCH2CH2),,_,, (OCH2CH2),OH
`
`
`
`-> C11]-l23CO(OR) (0Cl-l2CHg),,_, OH
`
`
`
`
`+ short chain degradation products
`Scheme VI
`
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`
`
`where R is the sorbitan ring, /1 is the number of oxyethylene groups
`
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`originally present, and x is the number of oxyethylene groups peroxidized
`
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`and subsequently degraded.
`
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`
`Hydrolysis of Polysorbate—Hydrolysis, which may be acid, base.
`
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`
`
`Journal of Phannaceutical Sciences / 1617
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`
`Vol. 67, No. 12, December 1978
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`
`relevant to the stability of other polysorbates. Some data
`
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`
`
`on the purity of polysorbates 40, 60, and 80 indicative of
`
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`
`the state of commercial samples also are reported.
`
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`
`For comparison with cetomacrogol, the peroxide con-
`
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`
`tent, pH, and total acidity were measured as parameters
`
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`
`of chemical decomposition; cloud point and surface tension
`
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`were used as criteria of physical changes bearing on both
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`the decomposition process and the possible failure of the
`
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`surfactant as a solubilizer or emulsifier in a formulation.
`
`THEORETICAL
`
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`Apart from peroxide formation, which might be expected to occur in
`
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`all oxyethylene-containing materials under suitable conditions,
`the
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`
`possibilit y of hydrolysis also must be considered in the polysorbate ester
`
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`surfactants. The first process is a chain reaction analogous to the per-
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`
`
`oxidation of oils and ethers (5) and occurs by autoxidation reactions as
`shown in Schemes l and II:
`
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`
`
`RH + light or catalyst ~> R‘ + H‘
`
`Scheme I—Initiation
`
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`
`
`R‘ + 02 —> R00‘
`
`
`
`R00‘ + RH A ROOH + R'
`
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`
`
`Scheme Il—Propugution
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`
`Free radicals also may be formed by the processes in Scheme lll and
`
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`removed by those in Scheme lV:
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`R00]-l
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`+ R00‘ + H‘
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`
`
`ROOH 4- R0‘ + ‘OH
`
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`
`
`ZROOH --> R00‘ + R0‘ + H20
`Scheme III
`
`
`(
`)
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`
`
`2R0§ —:> inactive products
`lb)
`
`
`
`
`
`R05 + R‘ —> inactive products
`1
`)
`
`
`2R’ —C-> inactive products
`Scheme IV—~Terminalion
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`On the basis that the peroxidation occurs in the hydrophilic chain in
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`pnlyoxyethylene surfactants (1, 2), the initiation and propagation steps
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`(Schemes I and II) yield hydroperoxides of the oxyethylene units ac-
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`cording to Scheme V:
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`—(0CH2CH2)0- + 02 ”* \(0CH2CH)0-
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`OOH
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`Scheme V
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`The hydroperoxide concentration is measured iodometrically in oils
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`and polyglycol solutions, and results are expressed as the hydroperoxide
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`equivalent to the iodine liberated in milliequivalents or millimoles per
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`liter (5, 6). Recalculation with respect to the weight of material under-
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`going autoxidation gives the peroxide number (P.N.), expressed in mil-
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`liequivalents per kilogram of surfactant (1, 5).
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`Peroxide Decomposition—Degradation of both hydroperoxides and
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`peroxide free radicals may occur by a number of routes (7-10), as sum-
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`marized elsewhere (1 1, 12). The peroxide formation rate during the initial
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`stage of propagation is normally faster than that of its decomposition.
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`The rates subsequently become equal, giving rise to a short plateau
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`representing a temporary steady state, following which decomposition
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`is the faster reaction. The analytical data give the residual peroxide
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`content, which is determined by the rates of the various simultaneous
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`reactions involved (8, 9); these reactions are temperature, concentration,
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`and catalyst dependent in aqueous surfactant systems (1).
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`The typical kinetic pattern, showing the rise to the maximum P.N.
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`value and then the fall, is a clear indication ofautoxidation with degra-
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`dation. The presence of more than minimal quantities of peroxide (P.N.
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`> 5) signifies that autoxidation is underway and has probably passed the
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`lag phase, but a low P.N. value may also be obtained if the decomposition
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`Page 2 of 6
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`.. in
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`HYDROLYSIS,%
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`/|( 00035
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`/,’
`Ii 0.0011]
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`ACIDITY,mEq/g
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`Table I~Peroxide Formation in Polysorbate 20 at Different
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`Temperatures
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`Time of“
`Id .
`“ uctm“ Maximum’ Maximum
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`I‘.N.,
`Period”,
`P.N.,
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`mEq/kg
`hr
`days
`<2
`50
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`5.5
`168
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`-
`75
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`<2
`130
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`10
`368
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`—
`268
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`17
`<5
`440
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`(>20)
`~50
`(>70)
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`(>20)
`—
`(>70)
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`(>25)
`<24
`(>140)
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`(>50)
`144
`(>70)
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`(>50)
`_
`(>50)
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`" C = catalyst (1 X 10"M CuSO4). L = light. and D = dark. ” Based on P.N.of
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`5 (1). ‘ Numbers in parentheses indicate values at end of experiment before ter~
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`minatlon was reached.
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`Temper-
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`70°
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`Conditions“
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`0O F‘+ F‘L‘
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`O
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`T‘
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`OUI"+UF‘+Ul“+U 7‘
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`or solvent catalyzed (15), would be expected to proceed at an increasing
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`rate as the pH falls on formation of acidic degradation products, following
`the reaction in Scheme VII:
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`C11H23CO(OR) (OCH-_;CHg),,0I-I
`" C1]H2;§CO0H +
`Scheme VII
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`Of the products, lauric acid is stable and micelle soluble. The micelle
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`saturation point would be determined by the increasing quantity of lauric
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`acid and decreasing quantity of micelles, both functions of the hydrolysis
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`rate, and also by changes in the micelle-solubilizing capacity as a result
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`of chain shortening. The sorbitan polyglycol would pass out of the mi-
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`cellar phase, raising the aqueous concentration of hydrophilic solute
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`progressively and possibly reducing the cloud point of the surfactant. It
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`would he expected to undergo peroxidation and degradation by the re-
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`actions outlined for the parent surfactant; but since the process would
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`occur outside the micellar phase, the rate constants and mechanism might
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`not be identical with those of the polysorbate.
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`Turbidity in the solution could be the result of either lauric acid sep-
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`aration or reduction of the cloud point.
`EXPERIMENTAL
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`Polysorbate 20 was neutralized before use to pH 6.00 with sodium
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`hydroxide.
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`Autoxidation was effected under the same conditions as used previ-
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`ously in studies on cetomacrogol (1, 2). The method was designed to en-
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`sure adequate agitation and free access of air during storage. without loss
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`of solvent by evaporation. Therefore, the oxygen concentration remained
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`constant and was not rate limiting.
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`The acid content, pl-I, cloud point (1), and surface tension (2) were
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`determined as described previously.
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`The peroxide number was determined using the spectrophotometric
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`method developed for determining hydroperoxide in micellar solutions
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`(6). Final readings were made at a concentration level of 1% polysorbate
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`containing pH 6.00 buffer and potassium iodide at the same concentra-
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`tions as described for cetomacrogol. Readings were taken at 360 nm (e
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`= 11,400) and were time independent in this system. Dilutions, when
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`necessary, were made with solutions of 1% polysorbate 20; readings were
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`corrected by deducting the amount of iodine liberated by the polysorbate
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`content of the diluent. The polysorbate used as the diluent was stored
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`under nitrogen and refrigerated.
`RESULTS AND DISCUSSION
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`Development of Peroxides—The three stages of autoxidation—uiz ,
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`induction. propagation, and termination, were observed (Fig. 1). The
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`peroxide number (P.N.) values at each stage varied with conditions but
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`were comparable with those of cetomacrogol under parallel conditions
`(1).
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`Elevation of temperature from 25 to 70° reduced the induction period
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`and raised the peroxide formation rate under all conditions. Copper
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`sulfate and light had the expected catalytic effects, shortening induction
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`and raising the peroxide formation rate relative to the dark uncatalyzed
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`reaction. Pronounced catalysis of peroxide breakdown by the metal ions
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`at 60 and 70° occurred (1, 3, 11), with a shorter time required to reach the
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`1678 I Journal of Pharmaceutical Sciences
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`Vol. 67, No. 12, December 1978
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`DAYS
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`Figure 2—— Rate of acid formation in 3% aqueous polysorbate 20 at 70,
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`60, 40, and 25° in daylight with no catalyst. Broken lines represent lauric
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`acid produced by hydrolysis and were calculated using the indicated
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`values of k (hour‘l), the first-order hydrolysis constant, on both acidity
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`and percent hydrolysis scales.
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`maximum P.N. value and a lower value of P.N. obtained (Table I). Ele-
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`vation of temperature also catalyzed the decomposition of peroxide.
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`Generally speaking, the greater the peroxide decomposition rate, the
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`lower was the P.N. value at the termination stage, in which the degra-
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`dation rate of the peroxides equaled or exceeded their formation rate.
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`The enhancement of formation and decomposition rates by the tem-
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`perature—catalyst combinations was such that termination was reached
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`even at 40°; yet at this temperature, the rate balance brought about the
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`highest P.N. value. Even at 25° with the catalyst, the PN. rose to 150
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`within 25 days, and there was virtually no induction period (Fig. 1c).
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`pH and Acidity—As observed with cetomacrogol (1), the increase in
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`acidity (Fig. 2) continued after the PN. fell and was, therefore, the most
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`reliable factor for following the degree of deterioration. The phenomenon
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`of incessant increase in acidity in short chain polyglycols also was de-
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`scribed by McKenzie (11). The rate of development of acidity had an
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`inverse relation to the time of onset of propagation (Fig. 1:1) and also to
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`the initial rate of formation of peroxides.
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`The pH value approached 4.0 at 25 and 40°, whereas at 60 and 70° it
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`fell rapidly and continuously to 2.5 (Fig. 3a). The relation between pH
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`and acid concentration, c (Fig. Rb), indicated that the acids developed
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`at the lower temperatures contained weaker functions. Log acid con-
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`centration—pI-I plots were linear at 40, 60, and 70° and tended to converge
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`at high acidity. By use of the equation pH = 1/2 (pKa + log c), the inter-
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`cepts at log 0 = 0 enabled estimation of apparent pKa values‘, which were
`3.6 i 0.5 at 60 and 70° and 4.9 d: 0.8 at 40°.
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`These results suggest the presence of a larger fraction of stronger acids
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`at the higher temperatures, which is consistent with greater rupture of
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`the oxyethylene chains (Scheme VI). Indeed, in cetomacrogol, an ether
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`for which hydrolysis is not to be expected, formic acid (pKa 3.75) con-
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`stituted 50% or more of the acid formed during the initial stages of aut-
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`oxidation under drastic conditions (12). The weaker acids present at 25°
`could be constituted of micelle-solubilized lauric acid (Scheme VII),
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`carboxylated surfactant, or acetic acid (12).
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`The hydrolysis rate of polysorbate 80 was reported to be relatively
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`constant and lowest between pH 3 and 7.6, increasing rapidly as a func-
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`tion of pH below 3 and above 7.6 (15). There was little difference between
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`the hydrolysis rates of different polysorbate esters. To estimate the degree
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`‘ These pKa values represent mixed acid systems. The linearity over one order
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`of concentration could indicate that the acid mixture has a relatively constant
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`composition over this region in the aqueous phase. Lyophobic acids solubilized in
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`the micelles would have a lesser influence on the experimental pH (16), and the pKa
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`estimate would relate to their apparent pKa values in the micellar solutions and
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`not to their aqueous pKa values.
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`Page 3 of 6
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`73%,dvnes/cm
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`(Jo
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`EETOMAERUBUL 50°
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`0
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`10
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`DAYS
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`Figure 4—Rate of change of surface tension, 7, of 3% aqueous poly-
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`sorbate 20 at 70, 60, 40, and 25° in daylight with no catalyst. (Cetoma~
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`crogol, 3% aqueous solution at 50°, is shown for comparison.)
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`20
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`30
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`0
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`20
`0.5
`1.5
`1.0
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`DAYS
`ACIDITY, mEq/g
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`Figure 3—The pH change '0; 3% aqueous polysorbate 20 at 70, 60, 40,
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`and 25° in daylight with no catalyst, with time (a), and as a function
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`of total acidity at the corresponding time (b).
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`of hydrolysis expected at the temperatures used in the present work,
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`approximate rate constants, 12, were calculated based on the rate constant
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`reported for 0.02% polysorbate 80 at pH 3.95 and 80°, utilizing the re-
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`ported energy of activation (15). This calculation gave 12 values of 3.80,
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`2.65, 1.21, and 0.62 X 10-3 hr“1 at 70, 60, 40, and 25°, respectively, and
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`these values were used to calculate the respective quantities of lauric acid
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`that would be yielded at various times.
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`Some of these results have been included in Fig. 2 for comparison. The
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`values are of the same order as the observed acidities at 25 and 40°, al-
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`though somewhat nverestimatedz; the curves are also similar in form.
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`However, at 70 and 60°, the quantity of acid formed on storage after 3
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`and 6 days, respectively, greatly exceeded that expected theoretically
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`from hydrolysis. Again, the upward curvature indicates a rate rising with
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`time, as in cetomacrogol (1), characteristic of degradation processes of
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`the chain—reaction type3.
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`Surface Tension (7) Changes—The surface tension above the CMC
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`decreased on storage and ultimately reached a constant value (Fig. 4).
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`The rate of fall of 7 increased systematically with a temperature rise; the
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`minimum values for 739., developed at 60 and 70° were lower than at 40
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`and 25° (about 28 and 30 dynes/cm, respectively). The rate was much
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`greater than observed under parallel conditions in cetomacrogol (1); a
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`typical result is included in Fig. 4. There was a rank correlation among
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`the rates of fall of 7, increase of acidity, and fall in pH suggestive of an
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`acidic reaction product increasing in quantity with temperature and in-
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`fluencing surface tension.
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`For the consideration of the hydrolysis of the fatty acid ester, tem~
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`perature acceleration has already been discussed. The increasing quantity
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`9 Hydrolysis rates are concentration dependent in acid solution, with the rate
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`constant falling by some 60% at high surfactant concentration (15). This fact is the
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`etter when the estimated 1: values are reduced by some 60%, as do the acidities
`grobable reason for the discrepancy; the 40 and 25° experimental acidities accord
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`at 60 and 70“ during the first 2 days (Fig. 2). A similar reduction of the rate constant
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`at a high surfactant concentration also was observed for peroxide formation in ne-
`tomacrogol (1).
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`“ The general picture and conclusions drawn would not be altered by the errors
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`inherent in this treatment due to the approximated k values or to the possibility
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`that the hydrolysis data on which they are based (15) are uncorrected for acids
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`formed by autoxidation, which oould be a serious source of error at 80°, particularly
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`at low pH values.
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`of hydrophobic lauric acid formed would be largely solubilized in the
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`micelles. However, being in equilibrium with the surface, some acid would
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`tend to be adsorbed there, forming a mixed surface film with the more
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`hydrophilic surface~active monomers of the polysorbate. Mixed films are
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`closer packed and are expected to give lower surface tension values than
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`the separate amphiphiles (13). Confirmation that this was the probable
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`explanation was obtained by measurement of the surface tensions of
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`mixtures of lauric acid and polysorbate 20 (Fig. 5).
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`The 71% value of the polysorbate fell steeply and almost linearly as the
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`lauric acid concentration was raised to about 2.2% (w/w) in polysorbate
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`[A7/Ac: ~ -3.1 dynes/cm/1% (w/w) lauric acid in polysorbate or -61
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`dynes/cm/mEq of lauric acid/g of polysorbate). Above 2.2% lauric acid
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`(equivalent to about 10% hydrolysis of the polysorbate), the rate of fall
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`decreased greatly to 0.32 dyne/cm/1% but remained approximately linear
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`up to saturation concentration [9% (w/w) lauric acid, 0.45 mEq/g,
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`A25°
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`040°
`0
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`I:i——D 60°
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`0-:0 70 °
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`73%,dynes/cm
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`1.5
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`1.0
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`ACIDITY, mEq/g
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`Figure 5—Sur/‘ace tension o/3% aqueous palysorbate 20 at 70, 60, 40,
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`and 25° as a function of total acidity at the corresponding time. Broken
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`lines represent surface tension of lauric acid solutions in polysorbate
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`20 (1% w/v), expressed as milliequivalents of lauric acid per gram of
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`polysorbate (V); and experimental surface tensions at 7 ‘’ ‘‘corrected’'
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`for acidity due to autoxidation by plotting against hydrolyzed poly~
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`sorbate at the corresponding time using k = 0.00211 hr‘1 (0).
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`Journal of Pharmaceutical sciencesl 1679
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`Vol. 67, No. 12, December 1978
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`Page 4 of 6
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`This hypothesis was borne out by the 7—total acid plot (Fig. 5). At the
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`higher temperatures, the acid content was greatly in excess of the amount
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`of lauric acid that could be solubilized in