Buffers, buffering agents, and ionic equilibria

INTROD UCTION It is well known that many dru gs are unstabl e when exposed to certa in acidi c or ba sic conditio ns, an d such informat ion is rou tinely gather ed during the prefor mu lation stage of developm ent. When such inst abiliti es are identifie d, one tool of the form ulation scienc es is to include a buffering agent (or agents ) in the dosage form with the hope that such excipie nts will impar t suf ficient stabili ty to en able the formulat ion. The proper ties that enable buffering agents to functi on as such is derived from their qualities as weak acids or bases, and have their ro ots in their respect ive ionic equilibria . AUTOIO NIZATIO N OF WA TER Even the purest grade of water contain s low concentra tions of ions that can be detect ed by means of ap pro priate cond uctivity measu rements. These ions aris e from the transfer of a proton from a water molecule to another: H2O þ H2O H3O þ þ OH
ð1Þ In Eq. (1), H3Oþ is known as the hydronium ion, and OH
is known as the hydroxide ion. This reaction is reversible, and the reactants are known to proceed only slightly on to the products. Approximating the activity of the various species by their concentrations, one can write the equilibrium constant for this reaction as KC ¼ ½H3Oþ½OH
 ½H2O 2 ð2Þ In aqueous solutions, the concentration of water is effectively a constant (55.55 M), and so Eq. (2) simpli fies to: KW ¼ ½H3Oþ½OH
 ð3Þ KW is known as the autoionization constant of water, and is sometimes identified as the ion product of water. The magnitude of KW is very small, being equal to 1.007  10
14 at a temperature of 25C.1 For the sake of convenience, Sørensen proposed the ‘‘p’’ scale, where numbers such as KW would be

Buffers, Buffering Agents, and Ionic Equilibria

Harry G Brittain

Center for Pharmaceutical Physics, Milford, New Jersey, U.S.A

INTRODUCTION

It is well known that many drugs are unstable when

exposed to certain acidic or basic conditions, and such

information is routinely gathered during the

preformu-lation stage of development When such instabilities

are identified, one tool of the formulation sciences is

to include a buffering agent (or agents) in the dosage

form with the hope that such excipients will impart

suf-ficient stability to enable the formulation The

proper-ties that enable buffering agents to function as such is

derived from their qualities as weak acids or bases, and

have their roots in their respective ionic equilibria

AUTOIONIZATION OF WATER

Even the purest grade of water contains low

concentra-tions of ions that can be detected by means of

appro-priate conductivity measurements These ions arise

from the transfer of a proton from a water molecule

to another:

In Eq (1), H3Oþ is known as the hydronium ion,

and OH
is known as the hydroxide ion This reaction

is reversible, and the reactants are known to proceed only

slightly on to the products Approximating the activity of

the various species by their concentrations, one can write

the equilibrium constant for this reaction as

KC ¼ ½H3O

þ½OH


In aqueous solutions, the concentration of water is

effectively a constant (55.55 M), and so Eq (2)

simpli-fies to:

KW is known as the autoionization constant of

water, and is sometimes identified as the ion product

of water The magnitude of KW is very small, being

equal to 1.007  10
14at a temperature of 25C.[1]

For the sake of convenience, Sørensen proposed the

‘‘p’’ scale, where numbers such as KW would be

expressed as the negative of their base10 logarithms The value of pKWwould then be calculated as

and would have a value equal to 13.997 at 25C Defining pH as

and

then Eq (3) can then be expressed as

The autoionization of water is an endothermic reac-tion, so KWincreases as the temperature is increased.[1] This temperature dependence is plotted in Fig 1

IONIC EQUILIBRIA OF ACIDIC AND BASIC SUBSTANCES

Of the numerous definitions of acids and bases that have been employed over the years, the 1923 defini-tions of J N Brønsted and T M Lowry have proven

to be the most useful for discussions of ionic equilibria

in aqueous systems According to the Brønsted–Lowry model, an acid is a substance capable of donating a proton to another substance, such as water:

The acidic substance (HA) that originally donated the proton becomes the conjugate base (A
) of that substance, because the conjugate base could conceiva-bly accept a proton from an even stronger acid than the original substance One can write the equilibrium constant expression corresponding to Eq (8) as

KC ¼ ½H3O

þ½A


Bio-V–Buffer But because [H

2O] is a constant, one can collect the constants on the left-hand side of the equation to

derive the acid ionization constant expression:

KA ¼ ½H3O

þ½A


And, of course, one can define pKAas

A strong acid is a substance that reacts completely

with water, so that the acid ionization constant defined

in Eq (10) or (11) is effectively infinite This situation

can only be achieved if the conjugate base of the strong

acid is very weak A weak acid will be characterized by

an acid ionization constant that is considerably less

than unity, so that the position of equilibrium in the

reaction represented in Eq (8) favors the existence of

unreacted free acid

A discussion of the ionic equilibria associated with

basic substances parallels that just made for acidic

sub-stances A base is a substance capable of accepting a

proton donated by another substance, such as water:

The basic substance (B) that originally accepted the

proton becomes the conjugate acid (BHþ) of that

substance, because the conjugate acid could conceivably donate a proton to an even stronger base than the original substance The equilibrium constant expression corresponding to Eq (12) is:

KC ¼ ½BH

þ½OH


Because [H2O] is a constant, the constants are col-lected on the left-hand side of the equation to derive the base ionization constant expression:

KB ¼ ½BH

þ½OH


pKBis defined as

A strong base is a substance that reacts completely with water, so that the base ionization constant defined

in Eq (14) or (15) is effectively infinite This situation can only be realized if the conjugate acid of the strong base is very weak A weak base will be characterized by

a base ionization constant that is considerably less than unity, so that the position of equilibrium in the reaction represented in Eq (12) favors the existence

of unreacted free base

IONIC EQUILIBRIA OF CONJUGATE ACIDS AND BASES

Once formed, the conjugate base of an acidic substance (i.e., the anion of that acid) is also capable of reacting with water:

Because aqueous solutions of anions are commonly prepared by the dissolution of a salt containing that anion, reactions of the type described by Eq (16) are often termed hydrolysis reactions Eq (16) is necessa-rily characterized by its base ionization constant expression:

KB ¼ ½HA½OH



and a corresponding pKBdefined in the usual manner, but because

it follows that

KB ¼ ½HAKW

Temperature (°C)

Fig 1 Temperature dependence of the autoionization

con-stant of water (From Ref.[1].)

Eq (19) contains the right-hand side expression of

Eq (10), so one deduces that

or

The same relation between ionization constants of a

conjugate acid–base pair can be developed if one were

to begin with the conjugate acid of a basic substance,

so Eq 21 is recognized as a general property of

conju-gate acid–base pairs

IONIC EQUILIBRIA OF BUFFER SYSTEMS

A buffer can be defined as a solution that maintains an

approximately equal pH value even if small amounts of

acidic or basic substances are added To function in

this manner, a buffer solution will necessarily contain

either an acid and its conjugate base, or a base and

its conjugate acid

The action of a buffer system can be understood

through the use of a practical example Consider acetic

acid, for which KA ¼ 1.82  10
5(pK ¼ 4.74) The

following pH values can be calculated (for solutions

having a total acetate content of 1.0 M) using its acid

ionization constant expression:

When an acidic substance is added to a buffer

sys-tem it would immediately react with the basic

compo-nent, as a basic substance would react with the acidic

component One therefore concludes from the table

that the addition of either 0.1 M acid or 0.1 M base

to a buffer system consisting of 0.5 M acetic acid and

0.5 M acetate ion would cause the pH to change by

only 0.18 pH units This is to be contrasted with the

pH changes that would result from the addition of

0.1 M acid to water (i.e., 7.0 to 1.0, for a change of

6.0 pH units), or from the addition of 0.1 M base to

water (i.e., 13.0 to 1.0, also for a change of 6.0 pH

units)

A very useful expression for describing the

proper-ties of buffer system can be derived from consideration

of ionization constant expressions For an acidic substance, Eq (10) can be rearranged as

½H3Oþ ¼ KA½A



Taking the negative of the base 10 logarithms of the various quantities yields the relation known as the Henderson–Hasselbach equation:

Eq (23) indicates that when the concentration of acid and its conjugate base are equal (i.e., [HA] ¼ [A
]), then the pH of the solution will equal the pKAvalue Therefore, a buffer system is chosen so that the target

pH is approximately equal to the pKAvalue

Viewed in this light, a buffer system can be envi-sioned as a partially completed neutralization reaction

where comparable amounts of HA and A
are present

in the solution The buffer region within a neutraliza-tion reacneutraliza-tion is shown in Fig 2, where the horizontal region in the graph of anion concentration and

[acetate]

pH

Fig 2 Neutralization curve obtained during the titration

of 1.0 M acetic acid, plotted as a function of the acetate ion concentration

observed pH reveals the buffer region of the system

For practical purposes, the buffer region would extend

over [HA]/[A
] ratios of approximately 0.2 to 0.8

SELECTION OF AN APPROPRIATE

BUFFER SYSTEM

The selection of a buffer system for use in a

pharma-ceutical dosage form is relatively straightforward It

is evident from the preceding discussion that the most

important prerequisite for a buffer is the approximate

equality of the pKA value of the buffer with the

intended optimal pH value for the formulation

Knowledge of the pH stability profile of a drug

sub-stance enables one to deduce the pH range for which

formulation is desirable, and the basis for the most

appropriate buffer system would be the weak acid or

base whose pKA or pKB value was numerically equal

to the midpoint of the pH range of stability

There are, of course, other considerations that need

to be monitored, such as compatibility with the drug

substance Boylan[2] has provided a summary of the

selection criteria for buffering agents:

1 The buffer must have adequate capacity in the

desired pH range

2 The buffer must be biologically safe for the intended use

3 The buffer should have little or no deleterious effect on the stability of the final product

4 The buffer should permit acceptable flavoring and coloring of the product

A practical consequence of Eq (23) is that as long as the concentration of a buffer is not overcome by reac-tion demands, a buffer system will exhibit adequate capacity within 1 pH unit with respect to its pKAor

pKBvalue

The second criterion from the preceding list restricts buffering agents to those deemed to be pharmaceuti-cally acceptable A list of appropriate buffer systems

is provided in Table 1, along with values for their

pKAor pKB values sourced from the compilations of Martell and Smith.[3–6] The use of buffering agents is most critical for parenteral formulations, and it has been noted over the years that phosphate, citrate, and acetate are most commonly used for such pur-poses.[7,8] Ethanolamine and diethanolamine are also used to adjust pH and form their corresponding salts, whereas lysine and glycine are often used to buffer pro-tein and peptide formulations Akers[9] has reviewed the scope of drug–excipient interactions in parenteral formulations and has provided an overview of the effect of buffers on drug substance stability

Table 1 Acids and bases suitable for use as buffer systems in pharmaceutical products

Martell and Smith reference

BUFFERS IN PHARMACEUTICAL SYSTEMS

It is well known that the stability of many active

phar-maceutical substances can be strongly dependent on

the degree of acidity or basicity to which they are

exposed, and that a change in pH can cause significant

changes in the rate of degradation reactions For such

compounds, formulators commonly include a buffer

system to ensure the stability of the drug substance

either during the shelf life of the product, or during

the period associated with its administration

In addition, preformulation scientists routinely use

buffer systems to set the pH of a medium in which they

intend to perform experimentation For instance, the

pH stability profile of a drug substance is routinely

obtained through the use of buffers, and the pH

depen-dence of solubility is frequently measured using

buffered systems However, the possibility that the

buf-fer system itself may influence or alter the results must

be considered in these studies

Stabilization of Drug Substances

in Formulations by Buffers

As mentioned previously, the stability of parenteral

formulations is often established through the use of

buffer systems, and Table 2 contains a partial listing

of such systems.[7,8]

The inclusion of a phosphate buffer in homatropine

hydrobromide ophthalmic solution enabled

formula-tors to fix the solution pH at 6.8, enabling the product

to be lyophilized.[10]This lyophilized product could be

stored for extended periods without degradation

Tro-methamine was found to effect a stabilizing effect on

N-nitrosoureas (such as lomustine, carmustine, and

tauromustine) in aqueous solutions.[11]

It has been reported that replacing succinate buffer

with glycolate buffer improved the stability of

lyophi-lized g-interferon.[12] In this work, it was found that

the succinate buffer could crystallize in the frozen

state, which limited its ability to maintain the

appro-priate pH, and therefore led to degradation On the

other hand, use of the glycolate buffer appeared to

minimize the freeze-induced pH shifting, and the

lyophilized product exhibited superior solid-state

stability

However, the use of buffers in parenterals is not

always benign, and numerous instances have been

summarized where buffers or other excipients have

caused stability problems.[9] For instance, the

com-plexation of Ca(II) and Al(III) with phosphate buffer

solutions has been studied at great length, as well as

the kinetic characteristics of the subsequent

precipita-tion of calcium and aluminum phosphate salts.[13–17]

The use of metal complexing excipients, such as citric

acid or ethylenediaminetetraacetic acid, was found to

be useful in delaying the onset of precipitation

The use of buffering agents in solid dose forms is not as widespread as the use in parenteral products Nevertheless, the current Handbook of Pharmaceuti-cal Excipients lists calcium carbonate, monobasic and dibasic sodium phosphate, sodium and potassium citrates, and tribasic calcium phosphate as potential buffering agents.[18]

In one study, the effect of 11 different compounds representing various classes of buffering agents were studied with respect to their effect on the dissolution kinetics of aspirin from tablet formulations.[19] It was found that buffering agents capable of reacting with acidic substances to evolve carbon dioxide (sodium bicarbonate, magnesium carbonate, or cal-cium carbonate) yielded the fastest dissolution rates, and hence were deduced to be more useful as tablet excipients Less effective were water-soluble buffering agents (such as sodium ascorbate or sodium citrate), and least effective were water-insoluble buffering agents (such as magnesium oxide, magnesium trisili-cate, dihydroxyaluminum aminoacetate, or aluminum hydroxide)

In another study, the kinetics of aspirin, salicylic acid, and salicyluric acid were followed upon oral

Table 2 Some of the buffer systems used to stabilize various parenteral products

Basis for buffering system Product trade name

Ceredase Cerezyme Duracillin A.S

Fentenyl citrate and Droperidol

Zantac injection Pregnyl Prolastin Synthroid

Methergine injection Priscoline injection

administration of aspirin as either an unbuffered tablet

or two buffered solutions.[20]Significant differences in

the absorption rates were observed, with the solution

having 16 mEq of buffer being the fastest, the solution

having 34 mEq of buffer being intermediate, and the

unbuffered tablet being the slowest These studies

demonstrate that inclusion of a buffering agent in a

tablet formulation of an acid-sensitive compound will

lead to the generation of better dosage forms

Use of Buffers to Study the pH Stability

Profile of Drug Substances

The evaluation of the pH stability profile of a drug

substance is an essential task within the scope of

pre-formulation studies Knowing the pH conditions under

which a given compound will be stable is of vital

importance to the chemists seeking to develop methods

of synthesis, to analytical scientists seeking to develop

methods for analysis, and to formulators seeking to

develop a stable drug product Typically, the

prefor-mulation scientist will prepare solutions of the drug

substance in a variety of buffer systems, and will then

determine the amount of drug substance remaining

after a predefined storage period However, for the

information to be useful, the investigator will also need

to verify that the buffer itself does not have an effect

on the observed reactions

The hydrolysis kinetics of vidarabine-50-phosphate

were studied at a variety of pH values that enabled

the compound to exist as its protonated, neutral, and

monoionized form.[21]It was found that the hydrolysis

reaction followed first-order kinetics at the five pH

conditions tested, and that the buffer system used did

not influence the reaction rates The pH–rate profile

suggested that even though the compound was most

stabile over pH 9.0 to 9.5, the stability at pH 7.4

(i.e., physiological pH) was more than adequate for

development of a parenteral formulation

The degradation kinetics of phentolamine

hydro-chloride were studied over a pH range of 1.2 to 7.2

and in various glycol solutions.[22] The kinetics were

determined to be first order over all pH values studied,

and a consideration of the ionization constant of the

compound indicated that only the protonated form

of the compound had been studied At relatively low

acidities, a pH-independent region (pH 3.1–4.9) was

noted for the hydrolysis, and the kinetics were not

affected by the concentration of buffer used However,

the degradation reaction was found to proceed at a

much faster rate at a pH of 7.2, and a small dependence

of rate constant on the concentration of phosphate in

the buffer system was noted

Other examples where buffers were successfully used

to study the pH stability of drug substances (and where

little or no effect could be ascribed to the buffer system used) include the chemical stability of diisoxazolyl-naphthoquinone[23] and metronidazole[24] in aqueous solution In another detailed study, the effect of pH, buffer species, medium ionic strength, and temperature

on the stability of azetazolamide was studied.[25] There are probably as many instances where buffer catalysis exerts a strong influence on pH stability studies as where no such effect exists For instance, the kinetics associated with the acid/base hydrolysis

of ciclosidomine were found to be strongly affected

by the concentration of buffer used to set the solution

pH for each study.[26] However, because a linear rela-tionship was found between buffer concentration and observed first-order rate constant, the effect of pH on the degradation was assessed by extrapolating to zero buffer concentration This information was used to deduce the buffer-independent pH–rate profile

In another study on solutions of spironolactone, the concentration of buffer was found to exert a strong influence on the degradation rate constants.[27] At the same time, the ionic strength of the medium did not appear to affect the rate constants The decomposition pathway for aqueous solutions of batanopride hydro-chloride was found to depend on the pH of the med-ium used for the study, although the concentration of buffer was found to exert catalytic effects.[28]

To those beginning work in this field, the study reported by Zhou and Notari on the kinetics of ceftazidime degradation in aqueous solutions may be used as a study design template.[29] First-order rate constants were determined for the hydrolysis of this compound at several pH values and at several tem-peratures The kinetics were separated into buffer-independent and buffer-dependent contributions, and the temperature dependence in these was used to calculate the activation energy of the degradation via the Arrhenius equation Ceftazidime hydrolysis rate constants were calculated as a function of pH, tempera-ture, and buffer by combining the pH–rate expression with the buffer contributions calculated from the buffer catalytic constants and the temperature dependencies These equations and their parameter values were able

to calculate over 90% of the 104 experimentally deter-mined rate constants with errors less than 10%

Use of Buffers to Study the pH Dependence

of Drug Substance Solubility

An evaluation of the effect of pH on the aqueous solu-bility of a drug substance is an essential component

of preformulation research, and such work is usually conducted along with determinations of ionization constants, solubilization mechanisms, and dissolution rates.[30]Methods for the determination of the solubility

of pharmaceutical solids have been discussed at length,[31]

and a large number of pH–solubility profiles have been

published in the 30 volumes of the Analytical Profiles

series.[32–34] A general treatment of the characteristics

of the pH–solubility profiles of weak acids and bases

is available.[35]

When the pH conditions used for a given solubility

determination are set through the use of buffers, the

possible solubilization of the buffering systems must

be established For instance, no buffer effect was

reported during the determination of the solubilities

of trimethoprim and sulfamethoxazole at various pH

values.[36] On the other hand, correction for buffer

effects was made during studies of some

isoxazolyl-naphthoquinone derivatives.[37]

With the continuing development of compounds

exhibiting low degrees of intrinsic aqueous solubility,

the combination of pH control and complexing agents

in formulations has become important, and buffers

play an important role in many of these formulations

A theoretical analysis of the synergistic effect observed

in the combined systems has been developed and used

to explain the solubilization noted for flavopiridol.[38]

In a subsequent work, the solubilization of this

sub-stance by pH control combined with cosolvents,

surfactants, or complexing agents was investigated.[39]

The combined effect of pH and surfactants on the

dissolution of piroxicam has been reported.[40] In this

system, the dissolution rate and solubility of the drug

substance could be well estimated by a simple additive

model for the effect of pH and surfactant, where the

total dissolved concentration equaled the summation

of the amount of dissolved non-ionized substance, the

amount of dissolved ionized substance, and the amount

of substance solubilized in the surfactant micelles It was

suggested that the model developed in this work could

be useful in establishing an in vitro–in vivo correlation

for piroxicam

An equilibrium-based model was proposed to

char-acterize the drug–surfactant interactions observed in

the system consisting of furbiprofen and polysorbate

80 in solutions of different pH.[41]The model reflected

both interactions and interdependence among all

drug-containing species, namely, non-ionized drug in water,

ionized drug in water, non-ionized drug in micelles,

and ionized drug in micelles The mathematical

treat-ment also enabled modeling of the drug solubilization

in the pH–surfactant solutions without requiring the

use of inappropriate approximations It was found that

the solubility data estimated by the proposed model

were more reliable when the surfactant concentration

was high in the system This finding confirmed that

that consideration of interrelations and interdependence

of all drug species in the various solutions was

appro-priate for this model

CONCLUSIONS Buffers and buffering agents have been widely used for the stabilization of pharmaceutical formulations, and this aspect has proven to be especially important for parenteral products Buffers and buffering agents have also been found to play a vitally important role during drug characterization studies, being vitally important

to the conduct of solubility and drug stability studies The range of pharmaceutically acceptable buffer sys-tems spans all useful pH values, and it can be said that there is a buffer available for every intended purpose

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