242
`
`CHAPTER 17
`
`Example 2-What is the change of pH on adding 0.1 mol ofNaOH to
`1 L of buffer solution 0.1 Min acetic acid and 0.1 Min sodium acetate?
`
`DETERMINATION OF PH
`
`(a) The pH of the buffer solution before adding NaOH is
`
`Colorimetry
`
`A relatively simple and inexpensive method for determining
`the approximate pH of a solution depends on the fact that some
`conjugate acid-base pairs (indicators) possess one color in the
`acid form and another color in the base form. Assume that the
`acid form of a particular indicator is red, and the base form is
`yellow. The color of a solution of this indicator will range from
`red when it is sufficiently acid, to yellow when it is sufficiently
`alkaline.
`In the intermediate pH range (the transition interval) the
`color will be a blend of red and yellow depending upon the ratio
`of the base to the acid form. In general, although there are
`slight differences between indicators, color changes apparent to
`the eye cannot be discerned when the ratio of base to acid form,
`or acid to base form exceeds 10:1. The use of Equation 83
`indicates that the transition range of most indicators is equal to
`the pKa of the indicator ± 1 pH unit, or a useful range of
`approximately two pH units. Standard indicator solutions can
`be made at known pH values within the transition range of the
`indicator, and the pH of an unknown solution can be deter(cid:173)
`mined by adding the indicator to it and comparing the resulting
`color with the standard solutions.
`Another method for using these indicators is to apply them
`to thin strips of filter paper. A drop of the unknown solution is
`placed on a piece of the indicator paper and the resulting color
`is compared to a color chart supplied with the indicator paper.
`These papers are available in a wide variety of pH ranges.
`
`Potentiometry
`
`Electrometric methods for the determination of pH are based
`on the fact that the difference of electrical potential between
`two suitable electrodes dipping into a solution containing hy(cid:173)
`dronium ions depends on the concentration (or activity) of the
`latter. The development of a potential difference is not a spe(cid:173)
`cific property of hydronium ions. A solution of any ion will
`develop a potential proportional to the concentration of that ion
`if a suitable pair of electrodes is placed in the solution.
`The relationship between the potential difference and con(cid:173)
`centration of an ion in equilibrium with the electrodes may be
`derived as follows. When a metal is immersed into a solution of
`one of its salts, there is a tendency for the metal to go into
`solution in the form of ions. This tendency is known as the
`solution pressure of the metal and is comparable to the ten(cid:173)
`dency of sugar molecules (eg, to dissolve in water). The metallic
`ions in solution tend, on the other hand, to become discharged
`by forming atoms, this effect being proportional to the osmotic
`pressure of the ions.
`For an atom of a metal to go into solution as a positive ion,
`electrons, equal in number to the charge on the ion, must be left
`behind on the metal electrode with the result that the latter
`becomes negatively charged. The positively charged ions in
`solution, however, may become discharged as atoms by taking
`up electrons from the metal electrode. Depending on which
`effect predominates, the electrical charge on the electrode will
`be either positive or negative and may be expressed quantita(cid:173)
`tively by the following equation proposed by Nernst in 1889:
`
`RT p
`E=nFlnP
`
`(84)
`
`where E is the potential difference or electromotive force, R is
`the gas constant (8.316 joules), Tis the absolute temperature,
`n is the valence of the ion, Fis the Faraday of electricity (96,500
`
`[base]
`pH = log [acid] + pKa
`
`0.1
`=logo.I+ 4.76 = 4.76
`
`(b) On adding 0.01 mo! ofNaOH per liter to this buffer solution, 0.01
`mol of acetic acid is converted to 0.01 mol of sodium acetate, thereby
`decreasing the concentration of acid to 0.09 M and increasing the
`concentration of base to 0.11 M . The pH is calculated as
`
`0.1 l
`9 + 4.76
`0
`= 0.087 + 4.76 = 4.85
`
`_0
`
`pH= log
`
`The change of pH in this case is only 0.09 unit, about 1/10 the change
`in the preceding example. The buffer capacity is calculated as
`
`mols of NaOH added
`change of pH
`
`0.0 l
`= 0.09 = O. ll
`
`Thus, the buffer capacity of the acetic acid-sodium acetate buffer solu(cid:173)
`tion is approximately 10 times that of the acetic acid solution.
`
`As is in part evident from these examples, and may be
`further evidenced by calculations of pH changes in other sys(cid:173)
`tems, the degree of buffer action and, therefore, the buffer
`capacity, depend on the kind and concentration of the buffer
`components, the pH region involved and the kind of acid or
`alkali added.
`STRONG ACIDS AND BASES AS "BUFFERS"-In the
`foregoing discussion, buffer action was attributed to systems of
`(1) weak acids and their conjugate bases, (2) weak bases and
`their conjugate acids, and (3) certain acid-base pairs that can
`function in the manner either of system 1 or 2.
`The ability to resist change in pH on adding acid or alkali is
`possessed also by relatively concentrated solutions of strong
`acids and strong bases. If to 1 L of pure water having a pH of
`7 is added 1 mL of 0.01 M hydrochloric acid, the pH is reduced
`to about 5. If the same volume of the acid is added to 1 L of
`0.001 M hydrochloric acid, which has a pH of about 3, the
`hydronium-ion concentration is increased only about 1 % and
`the pH is reduced hardly at all. The nature of this buffer action
`is quite different from that of the true buffer solutions. The
`very simple explanation is that when 1 mL of 0.01 M HCl,
`which represents 0.00001 g-eq of hydronium ions, is added to
`the 0.0000001 g-eq ofhydronium ions in 1 L of pure water, the
`hydronium-ion concentration is increased 100-fold (equivalent
`to two pH units), but when the same amount is added to the
`0.001 g-eq of hydronium ions in 1 L of 0.001 M HCl, the
`increase is only 1/100 the concentration already present.
`Similarly, if 1 mL of0.01 M NaOH is added to 1 L of pure wa(cid:173)
`ter, the pH is increased to 9, while if the same volume is added
`to 1 L of 0.001 molar NaOH, the pH is increased al(cid:173)
`most immeasurably.
`In general, solutions of strong acids of pH 3 or less, and
`solutions of strong bases of pH 11 or more, exhibit this kind of
`buffer action by virtue of the relatively high concentration of
`hydronium or hydroxyl ions present. The USP includes among
`its Standard Buffer Solutions a series of hydrochloric acid
`buffers, covering the pH range 1.2 to 2.2, which also contain
`potassium chloride. The salt does not participate in the buffer(cid:173)
`ing mechanism, as is the case with salts of weak acids; instead,
`it serves as a nonreactive constituent required to maintain the
`proper electrolyte environment of the solutions.
`
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`

`

`IONIC SOLUTIONS AND ELECTROLYTIC EQUILIBRIA
`
`RT
`I
`£ a. -
`ln(cid:173)
`nF
`c1
`
`or in terms of Briggsian logarithms
`
`I
`RT
`E = 2.30:3 nFlog, 0 -
`c'2
`
`243
`
`(88)
`
`(89)
`
`If for log10 1/c:! there is substituted its equivalent pH, the
`equation becomes
`
`RT
`E = 2.:30:3 nF pH
`
`(90)
`
`and finally by substituting numerical values for R, n, T, and F,
`and assuming the temperature to be 20°, the following simple
`relationship is derived:
`
`(91)
`
`E
`581
`
`_0
`
`0
`
`E = 0.0581 pH or pH=
`
`The hydrogen electrode dipping into a solution of known
`hydronium-ion concentration, called the reference electrode,
`may be replaced by a calomel electrode, one type of which is
`shown in Figure 17-4. The elements of a calomel electrode are
`mercury and calomel in an aqueous solution of potassium chlo(cid:173)
`ride. The potential of this electrode is constant, regardless of
`the hydronium-ion concentration of the solution into which it
`dips . The potential depends on the equilibrium that is set up
`between mercury and mercurous ions from the calomel, but
`the concentration of the latter is governed, according to the
`solubility-product principle. by the concentration of chloride
`ions, which are derived mainly from the potassium chloride in
`the solution. Therefore. the potential of this electrode varies
`with the concentration of potassium chloride in the electrolyte.
`Because the calomel electrode always indicates voltages
`that are higher, by a constant value. than those obtained when
`the normal hydrogen electrode chain shown in Figure 17-3 is
`used, it is necessary to subtract the potential due to the calomel
`electrode itself from the observed voltage. As the magnitude of
`this voltage depends on the concentration of potassium chloride
`in the calomel-electrode electrolyte, it is necessary to know the
`concentration of the former. For most purposes a saturated
`potassium chloride solution is used that produces potential
`
`+
`
`t·t;•f:C~•\:~i?~
`::~f{.:~\:~:_r~~~-
`
`-•, • •:. , , • . • , ~
`
`SOLUTION
`
`CALOMEL
`
`PURE MERCLRY
`
`Pt.A TN..t,A WIRE
`
`ORDINARY t,EROR,<
`
`Figure 17-4. Calomel electrode.
`
`p
`
`=il-1=-
`
`A
`
`------
`Figure 17-3. Hydrogen-ion concentration chain.
`
`8 =-=
`
`coulombs >. p is the osmotic pressure of the ions. and P is the
`solution pressure of the metal.
`Inasmuch as it is impossible to measure the potential dif(cid:173)
`ference between one electrode and a solution with any degree of
`certainty, it is customary to use two electrodes and to measure
`the potential difference between them. If two electrodes. both of
`the same metal. are immersed in separate solutions containing
`ions of that metal-at osmotic pressure p 1 and P:!, respectively
`-and are connected by means of a tube containing a nonreact(cid:173)
`ing salt solution (a so-called salt bridge). the potential devel(cid:173)
`oped across the two electrodes will be equal to the difference
`between the potential differences of the individual electrodes:
`thus.
`
`p , RT p.,
`RT
`E = E , - E-, = -
`In -
`-
`-
`In _.:c
`nF P ,
`nF P~
`-
`
`(85)
`
`As both electrodes are of the same metal, P 1
`equation may be simplified to
`
`P:! and the
`
`RT
`E ~ -
`In µ 1 -
`nF
`
`RT P ,
`RT
`In p ~ "" -
`In -:--
`-
`-n.F
`nF
`JJ ~
`
`(86)
`
`In place of osmotic pressures it is permissible, for dilute
`solutions, to substitute the concentrations c 1 and c'.! that were
`found (see Chapter 16). to be proportional to p 1 and p 2 • The
`equation then becomes
`
`RT
`E = -In (cid:173)
`nF
`-
`
`c,
`c~
`
`187 1
`
`If either c 1 or c:2 is known, it is obvious that the value of the
`other may be found if the potential difference. E. of this cell can
`be measured.
`For the determination of hydronium-ion concentration or
`pH. an electrode at which an equilibrium between hydrogen
`gas and hydronium ion can be established must be used in
`place of metallic electrodes. Such an electrode may be made by
`electrolytically coating a strip of platinum, or other noble
`metal, with platinum black and saturating the latter with
`pure hydrogen gas. This device functions as a hydrogen elec(cid:173)
`trode. Two such electrodes may be assembled as shown in
`Figure 17-3.
`In this diagram one electrode dips into Solution A, contain(cid:173)
`ing a known hydronium-ion concentration, and the other elec(cid:173)
`trode dips into Solution B, containing an unknown hydronium(cid:173)
`ion concentration. The two electrodes and solutions. sometimes
`called half-cells, then are connected by a bridge of neutral salt
`solution, which has no significant effect on the solutions it
`connects. The potential difference across the two electrodes is
`measured by means of a potentiometer, P. If the concentra(cid:173)
`tion, c1, of hydronium ion in Solution A is 1 N, Equation 87
`simplifies to
`
`Slayback Exhibit 1101, Page 22 of 39
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`
`

`

`244
`
`CHAPTER 17
`
`difference of 0.2488 V. Accordingly, before using Equation 86
`for the calculation of pH from the voltage of a cell made up of a
`calomel and a hydrogen electrode dipping into the solution to
`be tested, 0.2488 V must be subtracted from the observed
`potential difference. Expressed mathematically, Equation 92 is
`used for calculating pH from the potential difference of such a
`cell.
`
`ments available permit reading the pH directly and provide
`also for compensation of variations due to temperature in the
`range of 0° to 50° and to the small but variable asymmetry
`potential inherent in the glass electrode.
`
`E - 0.2488
`pH=
`0.0581
`
`PHARMACEUTICAL SIGNIFICANCE
`
`(92)
`
`•
`
`•
`
`In measuring the potential difference between the elec(cid:173)
`trodes, it is imperative that very little current be drawn from
`the cell, for with current flowing the voltage changes, owing to
`polarization effects at the electrode. Because of this it is not
`possible to make accurate measurements with a voltmeter that
`requires appreciable current to operate it. In its place a poten(cid:173)
`tiometer is used that does not draw a current from the cell
`being measured.
`There are many limitations to the use of the hydrogen
`electrode:
`It cannot be used in solutions containing strong oxidants such as
`•
`ferric iron, dichromates, nitric acid, peroxide, or chlorine or reduc(cid:173)
`tants such as sulfurous acid and hydrogen sulfide.
`It is affected by the presence of organic compounds that are reduced
`fairly easily.
`It cannot be used successfully in solutions containing cations that
`fall below hydrogen in the electrochemical series.
`• Erratic results are obtained in the measurement of unbuffered
`solutions unless special precautions are taken.
`It is troublesome to prepare and maintain.
`•
`As other electrodes more convenient to use now are available,
`the hydrogen electrode today is used rarely. Nevertheless, it is
`the ultimate standard for pH measurements.
`To avoid some of the difficulties with the hydrogen electrode,
`the quinhydrone electrode was introduced and was popular for
`a long time, particularly for measurements of acid solutions.
`The unusual feature of this electrode is that it consists of a
`piece of gold or platinum wire or foil dipping into the solution to
`be tested, in which has been dissolved a small quantity of
`quinhydrone. A calomel electrode may be used for reference,
`just as in determinations with the hydrogen electrode.
`Quinhydrone consists of an equimolecular mixture of qui(cid:173)
`none and hydroquinone; the relationship between these sub(cid:173)
`stances and hydrogen-ion concentration is
`
`Quinone + 2 Hydrogen ions + 2 Electrons ;:= Hydroquinone
`
`In a solution containing hydrogen ions the potential of the
`quinhydrone electrode is related logarithmically to hydronium(cid:173)
`ion concentration if the ratio of the hydroquinone concentration
`to that of quinone is constant and practically equal to l. This
`ratio is maintained in an acid solution containing an excess of
`quinhydrone, and measurements may be made quickly and
`accurately; however, quinhydrone cannot be used in solutions
`more alkaline than pH 8.
`An electrode that, because of its simplicity of operation and
`freedom from contamination or change of the solution being
`tested, has replaced both the hydrogen and quinhydrone elec(cid:173)
`trodes is the glass electrode. It functions because when a thin
`membrane of a special composition of glass separates two so(cid:173)
`lutions of different pH, a potential difference develops across
`the membrane that depends on the pH of both solutions. If the
`pH of one of the solutions is known, the other may be calculated
`from the potential difference.
`In practice, the glass electrode usually consists of a bulb of
`the special glass fused to the end of a tube of ordinary glass.
`Inside the bulb is placed a solution of known pH, in contact
`with an internal silver-silver chloride or other electrode. This
`glass electrode and another reference electrode are immersed
`in the solution to be tested and the potential difference is
`measured. A potentiometer providing electronic amplification
`of the small current produced is employed. The modern instru-
`
`In the broad realm of knowledge concerning the preparation
`and action of drugs few, if any, variables are so important as
`pH. For the purpose of this presentation, four principal types of
`pH-dependence of drug systems will be discussed: solubility,
`stability, activity, and absorption.
`
`Drug Solubility
`
`If a salt, NaA, is added to water to give a concentration c., the
`following reactions occur:
`
`If the pH of the solution is lowered, more of the A- would be
`converted to the unionized acid, HA, in accordance with Le
`Chatelier's principle. Eventually, a pH will be obtained, below
`which the amount of HA formed exceeds its aqueous solubility,
`S0 , and the acid will precipitate from solution; this pH can be
`designated as pH.P. At this point, at which the amount of HA
`formed just equals S0 , a mass balance on the total amount of
`drug in solution yields
`
`(93)
`
`Replacing [A-J as a function of hydronium-ion concentration
`gives
`
`C, = So + [H o+ I + K
`a
`3
`
`p
`
`( 94 )
`
`where K0 is the ionization constant for the conjugate acid, HA,
`and [H30+]P refers to the hydronium-ion concentration above
`which precipitation will occur. This equation can be rearranged
`to give
`
`Taking logarithms gives
`
`(95)
`
`(96)
`
`Thus, the pH below which precipitation occurs is a function of
`the amount of salt added initially, the pK0 and the solubility of
`the free acid formed from the salt.
`The analogous equation for salts of weak bases and strong
`acids (such as pilocarpine hydrochloride, cocaine hydrochlo(cid:173)
`ride, or codeine phosphate) would be
`
`So
`pH!> = pKa + log C., _ So
`
`(97 )
`
`in which pK0 refers to the protonated form of the weak base.
`Example-Below what pH will free phenobarbital begin to precipi(cid:173)
`tate from a solution initially containing 1.3 g of sodium phenobarbital/
`100 mL at 25°? The molar solubility of phenobarbital is 0.0050 and its
`pK0 is 7.41. The molecular weight of sodium phenobarbital is 254.
`The molar concentration of salt initially added is
`
`Slayback Exhibit 1101, Page 23 of 39
`Slayback v. Eye Therapies - IPR2022-00142
`
`

`

`_
`1:3
`g/ L
`C, = mol wt= 254 = 0.0;,I M
`
`ing specific emulsion systems, and the effect of pH upon them,
`may be found in Chapter 21.
`
`IONIC SOLUTIONS AND ELECTROLYTIC EQUILIBRIA
`
`245
`
`Drug Activity
`
`Drugs that are weak acids or weak bases-and hence may exist
`in ionized or nonionized form (or a mixture of bothJ-may be
`active in one form but not in the other; often such drugs have an
`optimum pH range for maximum activity. Thus, mandelic acid,
`benzoic acid, or salicylic acid have pronounced antibacterial
`activity in nonionized form but have practically no such activity
`in ionized form. Accordingly, these substances require an acid
`environment to function effectively as antibacterial agents. For
`example, sodium benzoate is effective as a reservative in 4%
`concentration at pH 7, in 0.06 to 0.1% concentration at pH 3.5
`to 4, and in 0.02 to 0.03% concentration at pH 2.3 to 2.4. Other
`antibacterial agents are active principally. if not entirely, in
`cationic form. Included in this category are the acridines and
`quaternary ammonium compounds.
`
`Drug Absorption
`
`The degree of ionization and lipoid solubility of a drug are two
`important factors that determine the rate of absorption of
`drugs from the gastrointestinal tract, and indeed their passage
`through cellular membranes generally. Drugs that are weak
`organic acids or bases, and that in nonionized form are soluble
`in lipids, apparently are absorbed through cellular membranes
`by virtue of the lipoidal nature of the membranes. Completely
`ionized drugs, on the other hand, are absorbed poorly, if at all.
`Rates of absorption of a variety of drugs are related to their
`ionization constants and in many cases may be predicted quan(cid:173)
`titatively on the basis of this relationship. Thus, not only the
`degree of the acidic or basic character of a drug, but also
`consequently the pH of the physiological medium (eg, gastric or
`intestinal fluid, plasma, cerebrospinal fluid) in which a drug is
`dissolved or dispersed- because this pH determines the extent
`to which the drug will be converted to ionic or nonionic form(cid:173)
`become important parameters of drug absorption. Further in(cid:173)
`formation on drug absorption is given in Chapter 58.
`
`REFERENCES
`l. Benet LZ. Goyan JE. J Pharm Sci 1965: 54: 1179.
`2. Riegelman Set al. J Pharm Sci 1962; 51: 129.
`3. Niebergall P,J et al. J Pharm Sci 1972; 61: 232.
`
`BIBLIOGRAPHY
`Conway BE. Ionic Hydration in Chemistry and Biophysics. Amsterdam:
`Elsevier. 1980.
`Denbigh K. The Principles of Chemical Equilibrium, 4th ed. London:
`Cambridge University Press. 1981.
`Freist>r H. Fernando Q. Ionic Equilibria in .4.nalytical Chemistry. New
`York: Wiley. 1966.
`Harned HS, Owen BB. The Physical Chemistry of Electrolytic Solutions.
`New York, Reinhold. 1958.
`
`0.051 - 0.005
`pHr = 7.41 + log
`05
`0
`= 7.41 + o.96 = s.:37
`Example-Above what pH will free cocaine begin to precipitate from
`a solution initially containing 0.0294 mo! of cocaine hydrochloride per
`liter? The pKb of cocaine is 5.59, and its molar solubility is 5.60 ,; 10-:i_
`
`_0
`
`pK., = pK,. - pK,. = 14.00 - 5.59 = 8.4 I
`0.0056
`pH,,= S.4! + log 0.0'.l94 - 0.0056
`= 8.41 + (-0.6:3) = 7.78
`
`Drug Stability
`
`One of the most diversified and fruitful areas of study is the
`investigation of the effect of hydrogen-ion concentration on the
`stability or. in more general terms. the reactivity of pharma(cid:173)
`ceutical systems. The evidence for enhanced stability of sys(cid:173)
`tems when these are maintained within a narrow range of pH,
`as well as of progressively decreasing stability as the pH de(cid:173)
`parts from the optimum range, is abundant. Stability (or in(cid:173)
`stability) of a system may result from gain or loss of a proton
`(hydrogen ion) by a substrate molecule-often accompanied by
`an electronic rearrangement-that reduces (or increases) the
`reactivity of the molecule. Instability results when the sub(cid:173)
`stance desired to remain unchanged is converted to one or more
`other, unwanted. substances. In aqueous solution, instability
`may arise through the catalytic effect of acids or bases-the
`former by transferring a proton to the substrate molecule, the
`latter by accepting a proton.
`Specific illustrations of the effect of hydrogen-ion concentra(cid:173)
`tion on the stability of medicinals are myriad; only a few will be
`given here, these being chosen to show the importance of pH
`adjustment of solutions that require sterilization.
`Morphine solutions are not decomposed during a 60-min
`exposure at a temperature of 100° if the pH is less than 5.5;
`neutral and alkaline solutions, however, are highly unstable.
`Minimum hydrolytic decomposition of solutions of cocaine oc(cid:173)
`curs in the range of pH of 2 to 5; in one study a solution of
`cocaine hydrochloride, initially at a pH of 5.7, remained stable
`during 2 months (although the pH dropped to 4.2 in this timel,
`while another solution buffered to about pH 6 underwent ap(cid:173)
`proximately 30% hydrolysis in the same time. Similarly, solu(cid:173)
`tions of procaine hydrochloride containing some hydrochloric
`acid showed no appreciable decomposition; when dissolved in
`water alone, 5% of the procaine hydrochloride hydrolyzed,
`whereas when buffered to pH 6.5. from 19 to 35% underwent
`decomposition by hydrolysis. Solutions of thiamine hydrochlo(cid:173)
`ride may be sterilized by autoclaving without appreciable de(cid:173)
`composition if the pH is below 5; above this, thiamine hydro(cid:173)
`chloride is unstable.
`The stability of many disperse systems. and especially of
`certain emulsions, is often pH dependent. Information concern-
`
`Slayback Exhibit 1101, Page 24 of 39
`Slayback v. Eye Therapies - IPR2022-00142
`
`

`

`This material may be protected by Copyright law (Title 17 U.S. Code)
`
`Ophthalmic Preparations
`
`CHAPTER43
`
`Gerald Hecht, PhD
`Senior Director, Pharmaceutical Sciences
`Alcon Laboratories
`Fort Worth, TX 76101
`
`Ophthalmic preparations are sterile products essentially free
`from foreign particles, suitably compounded and packaged for
`instillation into the eye. Ophthalmic preparations include so(cid:173)
`lutions, suspensions, ointments, and solid dosage forms. The
`solutions and suspensions are, for the most part, aqueous.
`Ophthalmic ointments usually contain a white petrolatum(cid:173)
`mineral oil base.
`Ophthalmic preparations can be grouped broadly into two
`divisions of major significance to the pharmacist. These include
`single or multidose prescription products and the category de(cid:173)
`scribed as OTC or over-the-counter ophthalmic products. The
`latter group has been subjected to a searching review and
`analysis by a body of experts as a part of the Food and Drug
`Administration's (FDA) OTC Drug Review process.
`The single dominant factor characteristic of all ophthalmic
`products is the specification of sterility. Any product intended
`for use in the eye regardless of form, substance, or intent must
`be sterile. This requirement increases the similarity between
`ophthalmic and parenteral products; however the physiology
`of the human eye in many respects imposes more rigid
`formulation requirements . This is considered in the following
`discussion.
`Preparations intended for the treatment of eye disorders
`can be traced to antiquity. Egyptian papyri writings describe
`eye medications. The Greeks and Romans expanded such uses
`and gave us the term collyria . Collyria refers collectively to
`materials that were dissolved in water, milk, or egg white for
`use as eyedrops. In the Middle Ages collyria included mydriatic
`substances to dilate the pupils of milady's eyes for cosmetic
`purposes, thus the term belladonna, or beautiful lady.
`From the time of belladonna collyria, ophthalmic technology
`progressed at a pharmaceutical snail's pace well into modern
`times. It was not until after World War II that the concept of
`sterility became mandatory for ophthalmic solutions. Prior to
`World War II and continuing into the 1940s very few ophthal(cid:173)
`mic preparations were available commercially or were de(cid:173)
`scribed officially. The USP XIV, official in 1950, included only
`three ophthalmic preparations, and all three were ointments.
`Preparations to be used in the eye, either solutions or oint(cid:173)
`ments, invariably were compounded in the community or hos(cid:173)
`pital pharmacy and were intended for immediate (prescription)
`use. Such preparation and prompt use is reflected in the phar(cid:173)
`maceutical literature of the times. The stability of ophthalmic
`preparations is discussed in terms of days or a few months.
`One of the most important attributes of ophthalmic products
`is the requirement of sterility. Even that, however, is a sur(cid:173)
`prisingly recent event. The USP XV in 1955 was the first
`official compendium to include a sterility requirement for oph(cid:173)
`thalmic solutions. The FDA in 1953 adopted the position that a
`nonsterile ophthalmic solution was adulterated. Sterile oph(cid:173)
`thalmic products were, of course, available prior to the mid-
`1950s; however the legal requirement of sterility dates only
`from 1955.
`
`The sterility requirements for ophthalmic ointments ap(cid:173)
`peared first in the USP XVIII, Third Supplement (1972). Prior
`to that date there was no legal requirement for a sterile oph(cid:173)
`thalmic ointment. This probably was due to the difficulty (at
`that time) of testing for sterility in such nonaqueous systems
`and also the anticipated difficulties in sterilizing and maintain(cid:173)
`ing sterile conditions during the manufacture and filling of
`ointments on a large scale.
`
`ANATOMY AND PHYSIOLOGY OF THE EYE
`
`The human eye is a challenging subject for topical administra(cid:173)
`tion of drugs. The basis of this can be found in the anatomical
`arrangement of the surface tissues and in the permeability of
`the cornea. The protective operation of the eyelids and lacrimal
`system is such that there is rapid removal of material instilled
`into the eye, unless the material is suitably small in volume
`and chemically and physiologically compatible with surface
`tissues. Figures 43-1 1 and 43-2 1 include pertinent anatomy of
`the human eye.
`EYELIDS-The eyelids serve two purposes: mechanical
`protection of the globe and creation of an optimum milieu for
`the cornea. The eyelids are lubricated and kept fluid-filled by
`secretions of the lacrimal glands and specialized cells residing
`in the bulbar conjunctiva. The antechamber has the shape of a
`narrow cleft directly over the front of the eyeball, with pocket(cid:173)
`like extensions upward and downward. The pockets are called
`the superior and inferior fornices (vaults), and the entire space,
`the cul-de-sac. The elliptical opening between the eyelids is
`called the palpebral fissure.
`EYEBALL-The wall of the human eyeball (bulbus, globe)
`is composed of three concentric layers.
`1. The outer fibrous layer.
`2. A middle vascular layer-the uvea or uveal tract, consisting of the
`choroid, the ciliary body, and the iris.
`3. A nervous layer-the retina.
`The outer layer is tough, pliable, but only slightly stretchable.
`In its front portion-the portion facing the outside world-the
`fine structure of the outer layer is so regular and the water
`content so carefully adjusted that it acts as a clear, transparent
`window (the cornea). It is devoid of blood vessels. Over the
`remaining two-thirds the fibrous coat is opaque (the white of
`the eye) and is called the sclera. It contains the microcircula(cid:173)
`tion, which nourishes the tissues of this anterior segment, and
`is usually white except when irritated and vessel dilatation
`occurs.
`The eyeball houses an optical apparatus that causes in(cid:173)
`verted reduced images of the outside world to form on the
`retina, which is a thin translucent membrane. The optical
`apparatus consists, in sequence, of the precorneal film, the
`cornea, the aqueous humor, the pupil, the crystalline lens, the
`
`821
`
`Slayback Exhibit 1101, Page 25 of 39
`Slayback v. Eye Therapies - IPR2022-00142
`
`

`

`822
`
`CHAPTER 43
`
`\
`
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`CRYPTS OF b
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`
`6---,,,.,..,+-- LiNz°s OF
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`GLAN DS OF
`ZEIS
`(SEBACEOUS)
`
`GLANDS OF MOLL A
`
`(SWEAT)
`
`VITR EOUS
`HUMOR
`
`Figure 43-1. The eye: vertical section .1
`
`vitreous humor, and the retina. The aqueous and vitreous
`humors are layers of clear fluid or gel-like material interposed
`between the solid structures. The pupil, a round centric hole in
`a contractile membranous partition (called the iris), acts as the
`variable aperture of the system. The crystalline lens is a re(cid:173)
`fractive element with variable power controlled and supported
`by a muscle incorporated in the ciliary body. The choroid is the
`metabolic support for the retina.
`The optical function of the eye calls for stability of its di(cid:173)
`mensions, which is provided partly by the fibrous outer coat;
`more effective as a stabilizing factor is the intraocular pres(cid:173)
`sure, which exceeds the pressure prevailing in the surrounding
`tissues. This intraocular pressure is the result of a steady
`production of specific fluid, the aqueous humor, which origi(cid:173)
`nates from the ciliary processes and leaves the eye by an
`intricate system of outflow channels. The resistance encoun(cid:173)
`tered during this passage and the rate of aqueous production
`are the principal factors determining the level of the intraocu(cid:173)
`lar pressure. In addition to this hydromechanical function, the
`aqueous humor acts as a carrier of nutrients, substrates, and
`metabolites for the avascular tissues of the eye.
`The bones of the skull join to form an approximately
`pyramid-shaped housing for the eyeball, called the orbit.
`CONJUNCTIVA-The conjunctiva! membrane covers the
`outer surface of the white portion of the eye and the inner
`aspect of the eyelids. In most places it is attached loosely and
`thus permits free movement of the eyeball. This makes possible
`subconjunctival injections. Except for the cornea the conjunc(cid:173)
`tiva is the most exposed portion of the eye.
`LACRIMAL SYSTEM-The conjunctiva! and corneal sur(cid:173)
`faces are covered and lubricated by a film of fluid secreted by
`the conjunctiva! and lacrimal glands. The secretion of the lac(cid:173)
`rimal gland, the tears, is delivered through a number of fine
`ducts into the conjunctiva! fornix. The s