>HiIifllli‘lflIiillllllllllflllllfllli! 0000000004921 2
`
`
`
`' v{EB/[INGTON
`
`
`fihe‘ sl-iehce and
`Practice of Pharmacy
`203!) EDITIGi}!‘
`"
` ';-~ ”N"
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 001
`
`

`

`
`
`Remington: The
`Science and Practice
`
`of Pharmacy
`
`Apotex, Inc. (IPR2019-00400), EX. 1014, p. 002
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 002
`
`

`

`
`
`Remington: The
`Science and
`
`Practice
`of Pharmacy
`
`ALFONSO R GENNARO
`
`1'5 ”“5" of the Editorial Board
`
`
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 003
`
`

`

`Editor: Daniel Limmer
`Managing Editor: Matthew J. Hauber
`Marketing Manager: Anne Smith
`
`Lippincott Williams & Wilkins
`
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`
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`y any means, including photocopying, or utilized
`by any information storage a
`nd retrieval system Without written permission
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`
`The publisher is not responsible (as a matter of product liability, negligence or
`otherwise) for any injury resulting from any material contained herein. This
`publication contains information relating to general principles of medical care
`which should not be construed as specific instructions for individual patients.
`Manufacturers’ product information and package inserts should be reviewed for
`current information, including contraindications, dosages and precautions.
`
`Printed in the United States of America
`
`Entered according to Act of Congress, in the year 1885 by Joseph P Remington,
`in the Office of the Librarian of Congress, at Washington DC
`
`Copyright 1889, 1894, 1905, 1907, 1917, by Joseph P Remington
`
`Copyright 1926, 1936, by the Joseph P Remington Estate
`
`Copyright 1948, 1951, by the Philadelphia College of Pharmacy and Science
`
`Copyright 1956, 1960, 1965, 1970, 1975, 1980, 1985, 1990, 1995, by the Phila-
`delphia College of Pharmacy and Science
`
`Copyright 2000, by the University of the Sciences in Philadelphia
`
`All Rights Reserved
`Library of Congress Catalog Card Information is available
`ISBN 0-683—306472
`
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`material. If they have inadvertently overlooked any, they will be pleased to make
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`
`The use of stractaral formula-s from USAN and the USP Dictionary of Drag
`Names is by permission of The USP Convention. The Convention is not respon-
`sible for any inaccuracy contained herein.
`Notice—This text is not intended to represent, nor shall it be interpreted to be, the
`eqnivalent of or a substitute for the official United States Pharmacopeia {USP)
`and/or the National Formalary (NF). In the event of any difference or discrep-
`ancy between the current official USP or NF standards of strength, quality,
`parity, packaging and labeling for drags and representations ofthem herein, the
`context and efi‘ect of the official compendia shall prevail.
`
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`at (800) 638-3030 or fax orders to (301) 824-7390. International customers
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`0001020304
`12345678910
`
`Apotex, Inc. (IPR2019-00400), EX. 1014, p. 004
`
`
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 004
`
`

`

`. A treatise on the theory
`.
`Remington: The Science and Practice of Pharmacy .
`and practice of the pharmaceutical sciences, with essential
`information about pharmaceutical and medicinal agents; also, a
`guide to the professional responsibilities of the pharmacist as the
`drug information specialist at the health team .
`.
`. A textbools and
`reference worls for pharmacists, physicians, and other practitioners of
`the pharmaceutical and medical sciences.
`
`EDITORS
`
`Alfonso R Gennaro, Choir
`
`Nicholas G Popovich
`
`Ara H Der Marderosian
`
`Roger L Schnoare
`
`Glen R Hanson
`
`Joseph B Schwartz
`
`Thomas Medwicls
`
`H Steve White
`
`AUTHORS
`
`The 1 19 chapters of this edition of Remington were written by the
`
`editors, by members of the Editorial Board, and by the authors
`
`listed on pages viii to x.
`
`Managing Editor
`
`John E Hoover, BSc (Pharm)
`
`Editorial Assistant
`
`Bonnie Brigham Packer, RNC, BA
`
`Director
`
`Philip P Gerbino 1995—2000
`
`Twentieth Edition—2000
`
`Published in the i80th year of the
`PHILADELPHIA COLLEGE OF PHARMACY AND SCIENCE
`
`Apotex, Inc. (IPR2019-00400), EX. 1014, p. 005
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 005
`
`

`

`
`
`
`
`240
`
`CHAPTER 17
`
`When the simplest amino acid salt, glycine hydrochloride, is
`dissolved in water, it acts as a diprotic acid and ionizes as
`‘NH,CH,coon + 3,0 : 'Nnacnmoo“ + HSO'
`‘ancsmoo + H20 -.-: NHZCH,coo-- + H,0*
`The form, *NH30H2000“, is an ampholyte because it also can
`act as a weak base:
`' NHgCHgCOO- + H-gO = ‘NchHgCOOH + OH"
`This type of substance, which carries both a charged acidic and
`a charged basic moiety on the same molecule is termed a
`zwittcrion. Because the two charges balance each other, the
`molecule acts essentially as a neutral molecule. The pH at
`which the zwitterion concentration is maximum is known as
`the isoelectric point, which can be calculated from Equation 75.
`On the acid side of the isoelectric point, amino acids and
`proteins are cationic and incompatible with anionic materials
`such as the naturally occurring gums used as suspending
`andl'or emulsifying agents. On the alkaline side of the isoelec-
`tric point, amino acids and proteins are anionic and incompat-
`ible with cationic materials such as benzalkoniuin chloride.
`
`Salts of Weak Acids and Weak Bases
`When a salt such as ammonium acetate (which is derived from
`a weak acid and a weak base) is dissolved in water, it under-
`goes the following reactions:
`BirA- “a? BH* + A"
`
`BH+ + 11,0 .—_- B + H,50'
`
`A' + H20:liA + OH-
`
`The total PBE for this system is
`[H301 + [HA] = [OH] + [Bl
`Replacing [HA] and [B] as a function of [H30+], gives
`
`[H;,O‘] +
`lHiOWC“ = [OH-1 +
`Km“
`”[30+] + K“
`[H304'l + K;
`in which C, is the concentration of salt, Kl, is the ionization
`constant of the conjugate acid formed from the reaction be-
`tween A‘ and water, and K; is the ionization constant for the
`protonated base, BH". In general, [HaCV'], [OH"]. Km and KJ,
`usually are smaller than C, and the equation simplifies to
`{H301 = was:
`‘73)
`
`(76}
`
`{77)
`
`Example—Calculate the pH of a 0.01 M solution of ammonium
`acetate. The ammonium ion has a K,‘ equal to 5.7:: X 1 "
`represents K; in Equation 78. Acetic aci has a K“ of 1.75 X 10' 5, which
`represents K5, in Equation 78:___________—————--
`[11,01 = “.75 x 10"” x 5.75 x 10""
`= 1.00 x 10' 7
`
`pH = — log (1.00 x 10' 7) = 7.00
`
`All of the assumptions are valid.
`
`If an acid or base is added to water, the pH of the latter is
`changed markedly, for water has no ability to resist change of
`pH; it is completely devoid ofbuffer action. Even a very weak
`acid such as carbon dioxide changes the pH of water, decreas-
`ing it from 7 to 5.7 when the small concentration of carbon
`dioxide present in air is equilibrated with pure water. This
`extreme susceptibility ofdistilled water to a change ofpH upon
`adding very small amounts of acid or base is often of great
`concern in pharmaceutical operations. Solutions of neutral
`salts, such as sodium chloride, similarly lack ability to resist
`change of pH on adding acid or base; such solutions are called
`unbnffered.
`Characteristic of buffered solutions, which undergo small
`changes ofpH on addition of acid or base, is the presence either
`of a weak acid and a salt of the weak acid, or a weak base and
`a salt of the weak base. An example of the former system is
`acetic acid and sodium acetate; and of the latter, ammonium
`hydroxide and ammonium chloride. From the proton concept of
`acids and bases discussed earlier, it is apparent that such
`buffer action involves a conjugate acid—base pair in the solu-
`tion. It will be recalled that acetate ion is the conjugate base of
`acetic acid, and that ammonium ion is the conjugate acid of
`ammonia (the principal constituent ofwhat commonly is called
`ammonium hydroxide).
`The mechanism of action of the acetic acid—sodium acct-ate
`buffer pair is that the acid, which exists largely in molecular
`(nonionized) form, combines with hydroxyl ion that may be
`added to form acetate ion and water; thus,
`01-13000}! -L 01-1 —-n- 01-13000 + H30
`The acetate ion, which is a base, combines with the hydrogen
`(more exactly hydronium) ion that may be added to form es—
`sentially nonionized acetic acid and water, represented as
`CH3COO' + 1130' —> CHECOOH + 11—30
`As will be illustrated later by an example, the change of pH is
`slight as long as the amount of hydronium or hydroxyl ion
`added does not exceed the capacity of the buffer system to
`neutralize it.
`The ammonia—ammonium chloride pair functions as a
`buffer because the ammonia combines with hydronium ion that
`may be added to form ammonium ion and water; thus,
`NH; 1- 1130'
`-3- NH; + H20
`Ammonium ion, which is an acid, combines with added hy-
`droxyl ion to form ammonia and water, as
`NH,‘ + OH —> NH, + H20
`Again, the change of pH is slight if the amount of adder
`hydronium or hydroxyl ion is not in excess ofthe capacity ofth-
`system to neutralize it.
`Besides these two general types of buffers, a third appear
`to exist. This is the buffer system composed of two salts, a
`monobasic potassium phosphate, KHZPO,” and dibasic pots:
`sium phosphate, KQHPO4. This is not, however, a new type I
`buffer; it is actually a weak-acidfconjugate-base buffer in whic
`an ion, H2P04 , serves as the weak acid, and HPOf" is i'
`conjugate base. When hydroxyl ion is added to this buffer tl
`following reaction takes place:
`mm, + on amen? + H30
`
`and when hydronium ion is added,
`HPO,2 + 1130'
`-—> H2130, + H20
`It is apparent that the mechanism of action of this type
`buffer is essentially the same as that of the weak-acidicon
`gate-base buffer composed of acetic acid and sodium acetat
`CALCULATIONSflA buffer system composed of a con
`gate acid—base pair, NaA—HA (such as sodium acetate 2
`acetic acid), would have a PBE of
`
`/ B
`
`UFFERS
`The terms buffer, buffer solution, and buffered solution, when
`used with reference to hydrogen-ion concentration or pH, refer
`to the ability of a system, particularly an aqueous solution, to
`resist a change of pH on adding acid or alkali, or on dilution
`with a solvent.
`
`Apotex, Inc. (IPR2019-00400), EX. 1014, p. 006
`
`___—‘
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 006
`
`

`

`[H30‘] + [HA] = [OI-1‘] -*- [A']
`
`(793
`
`Replacing [HA] and [K] as a function of hydronium—ion con—
`centration gives
`
`[H301 +
`
`[1130105
`[1430*] + Kn
`
`= [OH'] +
`
`'KU 0“
`-——r
`[1130‘] + Kr.
`
`(8m
`
`where 0,, is the concentration of the salt, NaA, and Cu is the
`:sncentration of the weak acid, HA. This equation can be re-
`arranged to give
`
`(81!
`(Cu ' {H301 + {OH l)_
`, _
`[H30 1 - Kg (Cr; + [“301 _ [OH I)
`
`Ir. general, both C,, and Ch are much greater than [H.JO‘J,
`which is in turn much greater than [Oi-H and the equation
`simplifies to
`
`
`K'YC"
`[ll;;0‘] = ‘1
`(1,
`
`or. expressed in terms of pH, as
`
`
`Cl
`[fa—plug l
`C“
`
`pH:
`
`(32]
`
`[83}
`
`This equation generally is called the Henderson—Hasselbalch
`equation. It applies to all buffer systems formed from a single
`conjugate acid—base pair, regardless of the nature of the salts.
`For example, it applies equally well to the following buffer
`systems: ammonia—ammonium chloride, monosodium phos-
`phate— disodium phosphate, and phenobarbital—sodium pheno-
`barbital. In the ammonia—ammonium chloride system, ammo—
`nia is obviously the base and the ammonium ion is the acid (0,,
`equal to the concentration ofthe salt). In the phosphate system,
`monosodium phosphate is the acid and disodium phosphate is
`the base. For the phenobarbital buffer system, phenobarbital is
`the acid and the phenobarbital anion is the base (0,, equal to
`the concentration of sodium phenobarbital}.
`As an example of the application of this equation, the pH of
`a buffer solution containing acetic acid and sodium acetate,
`each in 0.1 M concentration, may be calculated. The K“ of acetic
`acid, as defined above, is 1.8 X 10‘“, at 25".
`Solution:
`First, the pK" of acetic acid is calculated:
`
`pKn = — log K, = — log 1.8 X 10"“
`
`= — log 1.8 — log 10'IS
`= - 0.26 —- 1—5} = +4.?4
`
`Substituting this value into Equation 83:
`0.1
`pH =10g a + 4.74 = +4.T-l
`
`The Henderson-Hasselbalch equation predicts that any so-
`lutions containing the same molar concentration of acetic acid
`as of sodium acetate will have the same pH. Thus, a solution of
`0.01 M concentration of each will have the same pH, 4.74, as
`one of 0.1 M concentration of each component. Actually, there
`will be some difference in the pH of the solutions, for the
`activity coefficient of the components varies with concentration.
`For most practical purposes, however, the approximate values
`of pH calculated by the equation are satisfactory. It should be
`pointed out that the buffer of higher concentration of each
`component will have a much greater capacity for neutralizing
`added acid or base and this point will be discussed further in
`the discussion of buffer capacity.
`The Henderson-Hasselbalch equation is useful also for cal-
`culating the ratio of molar concentrations of a buffer system
`required to produce a solution of specific pH. As an example,
`
`IONIC SOLUTIONS AND ELECTROLYTIC EQUILIBRIA
`
`241
`
`suppose that an acetic acid-sodium acetate buffer of pH 4.5
`must be prepared. What ratio of the bufi'er components should
`be used?
`
`Solution:
`Rearranging Equation 83, which is used to calculate the pH ofweak
`acid—salt type buflers, gives
`
`[base] _ H
`“K”
`"3 [arid] ‘ p
`= 4.5 — 4.76 = —o.2-1 = (9.76 — 10)
`
`|
`
`
`__-
`W .
`_
`_
`[base]
`[acid] = anulog of (9.14) — 10) = 0.3.15
`
`The interpretation of this result is that the proportion of
`sodium acetate to acetic acid should be 0.575 mol of the former
`to 1 mol of the latter to produce a pH of 4.5. A solution con-
`taining 0.0575 mol of sodium acetate and 0.1 mol of acetic acid
`per liter would meet this requirement, as would also one con-
`taining 0.00575 mol of sodium acetate and 0.01 mol of acetic
`acid per liter. The actual concentration selected would depend
`chiefly on the desired buffer capacity.
`BUFFER CAPACITY—Elie ability of a buffer solution to
`resist changes in pH upon addition of acid or alkali may be
`measured in terms of buffer capacity. In the preceding discus-
`sion of buffers, it has been seen that, in a general way, the
`concentration of acid in a weak~acidfconjugate-base buffer de»
`termines the capacity to “neutralize” added base, while the
`concentration of salt of the weak acid determines the capacity
`to neutralize added acid. Similarly, in a weak-baser'conjugate-
`acid buffer the concentration of the weak base establishes the
`buffer capacity toward added acid, while the concentration of
`the conjugate acid of the weak base determines the capacity
`toward added base. When the buffer is equimolar in the con-
`centrations of weak acid and conjugate base, or of weak base
`and conjugate acid, it has equal buffer capacity toward added
`strong acid or strong base.
`Van Slyke, the biochemist, introduced a quantitative ex-
`pression for evaluating bufi‘er capacity. This may be defined as
`the amount, in gram-equivalents (g-eq) per liter, of strong acid
`or strong base required to be added to a solution to change its
`pH by 1 unit; a solution has a buffer capacity of 1 when 1 L
`requires 1 g-eq of strong base or acid to change the pH 1 unit.
`{In practice, considerably smaller increments are measured,
`expressed as the ratio of acid or base added to the change of pH
`produced.) From this definition it is apparent that the smaller
`the pH change in a solution caused by the addition ofa specified
`quantity of acid or alkali, the greater the buffer capacity of the
`solution.
`The following examples illustrate certain basic principles
`and calculations concerning buffer action and buffer capacity.
`
`Example I—What is the change of pH on adding 0.01 mol of NaOH
`to 1 L of 0.10 M acetic acid?
`
`{a} Calculate the pH of a 0.10 molar solution of acetic acid:
`
`[H301 = \-K..(-‘.. = ~.1.75 x10 4 x10 x 10 I = 4.13 x 10-3
`pH = — log 4.13 x10”: 2.33
`
`(bl 011 adding 0.01 mol of NaOH to a liter of this solution. 0.01 mol of
`acetic acid is converted to 0.01 mol of sodium acetate, thereby
`decreasing Ca to 0.09 M, and Cb = 1.0 X 10 '2 M. Using the
`Henderson-Hasselbach equation gives
`
`_
`0.01
`pH = 4.70 + log 0% =-1.76 — 0.51:) -— 3.8]
`
`The pH change is, therefore, 1.43 unit. The buffer capacity as defined
`above is calculated to be
` =0.(Jll
`mole of NaOH added
`change in pH
`
`Apotex, Inc. (IPR2019-00400), EX. 1014, p. 007
`
`
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 007
`
`

`

`
`
`242
`
`CHAPTER 17
`
`Example 2—What is the change of pH on adding 0.1 mol of NaOH to
`1 L of bufi'er solution 0.1 M in acetic acid and 0.1 M in sodium acetate?
`
`DETERMINATION OF PH
`
`
`
`(a) The pH of the buffer solution before adding NaOH is
`
`pH = log
`
`
`[base]
`[acid]
`0.1
`
`+ pig,
`
`=103a + 4.76 = 4.76
`
`(b’ 011 adding 0.01 mol of NaOH 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
`
`176
`”‘11
`1
`H
`‘3 “$909+"
`= 0.087 + 4.76 = 11.85
`
`The change of pH in this case is only 0.09 unit, about 12'10 the change
`in the preceding example. The buffer capacity is calculated as
`
`male of NaOH added
`change of pH
`
`0.01 —0
`‘0.o9_ '11
`
`Thus, the bufl'er capacity of the acetic acid—sodium acetate bufler solun
`tion is approximately 10 times that of the acetic acid solution.
`
`Calorimetry
`
`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 pKL, 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-
`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.
`
`As is in part evident from these examples, and may be
`further evidenced by calculations of pH changes in other sys-
`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. Ifto 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 geq of hydronium ions, is added to
`the 0.0000001 g—eq of hydronium ions in 1 L of pure water, the
`hydronium-ion concentration is increased IOU-fold (equivalent
`to two pH units), but when the same amount is added to the
`0.001 g-eq of hydroniurn ions in 1 L of 0.001 M HCl, the
`increase is only 11100 the concentration already present.
`Similarly, if 1 mL of 0.01 M NaOH is added to 1 L of pure wa-
`ter, the pH is increased to 9, while if the same volume is added
`to 1 L of 0.001 molar NaOI—I,
`the pH is increased al-
`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-
`ing mechanism, as is the case with salts ofweak acids; instead,
`it serves as a nonreactive constituent required to maintain the
`proper electrolyte environment of the solutions.
`Apotcx, Inc. (IPR2019-00400), EX. 1014, p. 008
`
`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-
`droniuin ions depends on the concentration (or activity) of the
`latter. The development of a potential difference is not a spe—
`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“
`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-
`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-
`tively by the following equation proposed by Nernst in 1889:
`
`E=R__Tln£
`nFP
`
`(84}
`
`where E is the potential difference or electromotive force, R is
`the gas constant (8.316 joules], T is the absolute temperature,
`it is the valence of the ion, F is the Faraday of electricity (96,500
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 008
`
`

`

`
`
`1264
`
`CHAPTER 67
`
`to tissue, decorporation, and diminished incorporation of calcium. 1'
`principal use is in the treatment of hgpercalcemia. It is not used alor
`as a source of potassium or phosphate in potassium or phospha'
`deficiency. It is a component of Potassium Phosphates, Potassium an
`Sodium Phosphates, and Dibosic Potassium and Sodium Phosphates.
`is also a reagent and pharmaceutical necessity for various buffers ar
`parenteral fluids. It is no longer used as a laxative; it may cau:
`diarrhea by the oral route. See Monobosic Potassium Phosphate fi
`other adverse effects.
`
`MONOBASIC POTASSIUM PHOSPHATE
`
`Potassium Phosphate Monobasic; K-Phos; Neutra-Phos
`[7778—77—0] K,HPO_, (136.091.
`Preparation—As for Dibosic Potassium Phosphate.
`Description—pH 15"}? aqueous solution: about 5.
`Solubility—l g in about 5 mL water.
`Comments—See Dihasic Potassium Phosphate for actions to d
`crease calcium absorption, depress calcium levels in plasma, and e'
`hance calcium excretion as pyrophosphate complex. The dibasic salt
`likewise used to treat hypercalcemta. It is used to treat uephrolithius
`when the stones are calcific. In this, the decrease in free calciu
`excretion into the urine decreases stone formation, and acidification
`the urine Ill-1.3130,
`causes acidosis! and £ree pyrophosphate ion fav:
`dissolution of stones. It is a component of Potassium Phosphates at
`Mouobosic Potassium and Sodium Phosphates and a pharmaceutic
`necessity for various parenteral fluids and buffers. Adverse effects a:
`diarrhea by the oral route lit is poorly absorbed orally and acts as r
`osmotic catharticl, hypocalcemia tparesthesias, confusion, weaknes
`muscle cramps, dyspnea, irregular heartbeat! when employed vigo
`ously in nonhypercalcemia patients. and the passing of loosened kidnl
`stones.
`
`POTASSIUM PHOSPHATES
`A mixture of monobasic and dibasic potassium phosphate in the rat
`described under each category below.
`Comments—For actions, uses, and adverse effects, see Dibas
`Potassium Phosphate and Monobo‘st'c Potassium Phosphate. Main
`used for hypercalcemia and hypophosphatemia.
`POTASSIUM AND SODIUM PHOSPHATES
`A mixture of mono- and dibasic potassium and sodium phosphates.
`Comments—See Dibosic Potassium. Phosphate, Monobosic Sodiu
`Phosphate, Dibasr‘c Sodium Phosphate (this pagel. The mixture is a
`vantageous in that it lessens the risk of sodium or potassium overlot
`from a single—entity preparation. It is used mainly for hypercolcem
`and hypophosphutemia.
`MONOBASIC POTASSIUM AND SODIUM
`PHOSPHATES
`A mixture of monobasic potassium and monobasic sodium phosphate
`Comments—See Monohosic Potassium Phosphate and Mrmobos
`Sodium Phosphate. The combination is used to acidify the urine for t]
`prevention and treatment of urolithiasis. The combination is advant
`geous in that it lessens the likelihood of excessive intake of eith
`sodium or potassium from that of the single-entity components.
`RINGER’S INJECTION—page 1248.
`LACTATED RINGER’S INJECTION—page 1248.
`SO DIU M ACETATE
`
`POTASSIUM GLUCONATE
`Kaon
`
`H i‘
`‘-0
`H
`'ClI-v—E—COOK
`0
`H
`
`HICI
`
`O—1
`.0
`{lg—-
`
`H0042—
`
`[299-27-4} C,;,H,,KOT £234.25).
`Preparation—Glucose may be oxidized to gluconic acid by various
`processes, eg, electrolytic oxidation of an alkaline solution, reaction
`with hypobromites, or fermentation using Aspergillus niger or other
`microorganisms. Neutralization with potassium hydroxide provides the
`salt.
`Desar'iption—W'hite to yellowish white, crystalline powder or
`granules; odorless; slightly bitter taste; stable in air; solutions slightly
`alkaline to litmus.
`Solubility—1 g in 3 mL water; practically insoluble in dehydrated
`alcohol, ether, or chloroform.
`Comments—A source ofpotassium for management of hypoholemic
`states, such as occur consequent to adrenocorticosteroid therapy or use
`ofthiazide diuretics, or for deliberate production of hyperkalemia, as for
`treatment of digitalis intoxication. The gluconate anion supposedly
`makes the compound better tolerated in the GI tract than is potassium
`chloride. It also is claimed that the potassium of the gluconate is
`absorbed high in the GI tract, above the location where mucossl lesions
`sometimes occur in combined thiazide-potassium therapy, whereas
`other salts are not absorbed so quickly. Such faulty suppositions and
`claims ignore the unavoidable chemical fact that irrespective of the salt
`used, potassium ion is only dissociable completely and hence is unaf-
`fected in its irritant actions and absorption by the anion in the
`compound.
`its sugar-coated tablets dissolve at a higher level than do enteric-
`coated tablets of potassium chloride but, by this very fact, are free to
`cause the irritation for which the chloride tablet was coated. The fact
`that it may cause nausea, vomiting, diarrhea, and abdominal discom-
`fort shows that the gluconate has no advantage over non-enteric-coated
`potassium chloride tablets. A full glass of water taken with either
`greatly reduces the irritant effects of either salt. Hypochloremia is a
`frequent accompaniment ofhy'pokalemia; in such instances the chloride
`definitely is preferred. Furthermore, since gluconate metabolizes to
`bicarbonate, it contributes to alkalosis, which also may be present in
`hypokalemia. Only in a hypokalemic, hyperchloremic acidosis {as in
`renal failure, dehydration, and occasional diabetic acidosis: is the drug
`rational; however, clinical experience indicates no obvious superiority
`over KC]. The use and toxicity of, and contraindications to, it are the
`same as for Potassium Chloride.
`
`POTASSIUM MIXTURES
`A number of potassium-containing products are mixtures of KCl and
`KHCOa; KCl, KHCO,,, and K100”; KCl, KHCO.3, and citric acid; K01,
`KHCO,,, and potassium citrate; ICE-1003 and citric acid; KC-l and potas-
`sium gluconate; KHCOQ, potassium citrate, and potassium acetate; and
`potassium citrate and potassium gluconate. Those that combine
`KHCOS with citric acid are efl‘ervescent; some effervescent preparations
`contain betaine.HCl or lysine.HCl in lieu of, or in addition to, citric acid.
`Those that are not reconstituted for effervescence are intended for their
`alkalinizing effects in addition to their effects to repair potassium
`deficits. KHCO3 and K2003 are directly all-calotic; potassium acetate,
`citrate, and gluconate all metabolize to KHCOS. Since hypokalemia
`usually is accompanied by othelosis, there are few situations in which
`an alkalinizing source of potassium is rational. Examples in which
`hypokalemia and acidosis coexist are renal failure, dehydration, and
`sometimes diabetic acidosis. Even in these, clinical experience is that
`KCl alone seems to be as useful as the combinations.
`
`DIBASIC POTASSIUM PHOSPHATE
`
`Sodium Acetate Trihydrate
`[6131-90-4] CBH,,r'aO.V,_.3HZO £136.08); anhydrous 1127-09-3] [82.03].
`Preparation—By neutralizing acetic acid with sodium carbonat
`Description—Colorless, transparent crystals or granular. crystt
`line powder; slightly bitter, saline taste; effloresces in warm, dry air; tl
`trihydrate liquefies at about 60°.
`Solubility—d g in 0.8 mL water or 19 mL alcohol.
`Comments—The acetate ion is metabolized rapidly and complete
`in the body; consequently, administration eventually is equivalent
`giving sodium bicarbonate. Solutions are stable and readily sterilize
`and this salt has been used for parenteral therapy of metabolic acidos
`and hyponatremia. It also may be used to alkalinize the urine. It is
`pharmaceutical necessity used in solutions for hemodialysis and pe]
`toneal dialysis.
`
`Potassium Phosphate Dibasic} Neutra-Phos
`[7758-11-4] KZHPO, :174.18).
`Preparation—By partial neutralization of phosphoric acid with
`potassium hydroxide or carbonate.
`Description—Granular powder; hygroscopic; pH (5% aqueous so-
`lution! about 8.5.
`Solubility—Very soluble in water.
`Comments—In the body, H130,2 anion interacts with calcium ion
`in a way that favors the deposition of both calcium and phosphate in
`bone salts and in other tissues depots. Some of the phosphate also is
`converted to pyrophosphate, which is a chelator of calcium, the calcium-
`Carbonic acid, monosodium salt; Baking Soda; Sodium
`pyrophosphate complex being excreted in the urine. Furthermore, high
`Acid Carbonate
`plasma phosphate levels decrease calcitriol levels and thus decrease
`Mon
`sodium rbonate 144-55-S] Nal-ICO.3 £84. 01!
`absorption of calcium. Thus KHPO, causes redAthbutmnfrom plasa
`potex I.nc “(iPR20i9-
`OZbO) x 1014 p 009
`
`SODIUM BICARBONATE
`
`Apotex, Inc. (IPR2019-00400), Ex. 1014, p. 009
`
`