Iniernatronal Jourrul of Pharmuceutrts, 33 (1986) 201-217 Elsevier 201 IJP 01121 Salt selection for basic drugs
`
`Philip L. Gould
`
`Pharmuceutical Research and Development Department, Pfizer Centrul Reseurch, Sandwich, Kent (U.K.) (Received 24 March 1986) (Accepted 30 May 1986) Key words: Salt form selection - Pharmaceutical salts
`
`Summary
`
`An attempt has been made using a Kepner-Tregoe decision analysis approach to provide rationale to salt selection for basic drugs. The selection objectives are reviewed in terms of the ‘essential’ (MUSTS) and ‘desirable’ (WANTS) issues. The desired characteristics of the salt form, given sufficient strength and toxicological suitability of the conjugate acid, are then discussed on the basis of the various pivotal physicochemical properties; melting point, aqueous solubility and dissolution rate, stability and hydrophobicity. Several trends are established which can then assist the decision of which range of salt forms to evaluate to overcome a particular problem with a basic drug. It is concluded that it is important to view the choice of salt form for development as a compromise, with particular focus on the correctly weighted desires to obtain the best balanced choice.
`
`Introduction
`
`and
`
`its conjugate acid. Thus, a clear compromise of properties for the salt form is required, but the difficulty remains of assessing which salt forms are best to screen for a particular drug candidate. Little, if any, literature has been devoted to discussing the compromise of properties for salt form selection. This review addresses the problem of salt form selection for basic drugs.
`
`Salt formation provides a means of altering the physicochemical and resultant biological char- acteristics of a drug without modifying its chem- ical structure. The importance of choosing the ‘correct’ salt form of a drug is well outlined in a published review (Berge et al., 1977) but, although salt form can have a dramatic influence on the overall properties of a drug, the selection of the salt form that exhibits the desired combination of properties remains a difficult semi-empirical choice. In making the selection of a range of potential salts, a chemical process group considers issues on the basis of yield, rate and quality of the crystalli- sation as well as cost and availability of the con- Correspondence: P.L. Gould, Pharmaceutical Group, Product Research and Development Laboratories, Cyanamid of Great Britain Limited, Gosport, Hams, U.K. jugate acid. The formulation and analytical groups are, on the other hand, concerned with the hygro- scopicity, stability, solubility and processability profile of the salt form, while the drug metabolism group is concerned with the pharmacokinetic aspects and the safety evaluation group on the toxicological effects of chronic and acute dosing of the drug
`
`Approach to the salt selection process
`
`Walking and Appino (1973) have used the Kepner-Tregoe (KT) techniques (Kepner and 0378-5173/86/$03.50 0 1986 Elsevier Science Publishers B.V. (Biomedical Division)
`
`Argentum EX1011
`
`Page 1
`
`

`

`“GO”/
`
`issues
`
`Organic:
`
`202 Tregoe, 1976) of decision analysis and potential problem analysis to aid the selection of a salt form. Although their application is more exem- plary of the KT method rather than of the specific application, the rational process decision analysis approach which defines essential and desirable attributes as ‘MUSTS’ and ‘WANTS, respec- tively, provides a route to initially address the problem of salt form selection.
`
`The major “GO”/“NO-GO” (MUSTS) issue for salt selection of an ionizable drug is the con- sideration of the relative basicity of the drug and the relative strength of the conjugate acid. Clearly to form a salt the pK, of the conjugate acid has to be less than or equal to the pK, of the basic centre of the drug. Thus the potential range of salts of drugs con- taining for example triazoyl bases (I; pK, - 2) is restricted to strong acids (mineral and sulphonic, but excluding the carboxylic), whereas imidazole bases (II; pK, 6-7) are far less restricted and the greatest scope for salt formation occurs for the aliphatic tertiary amines (III; pK, 9-10). t;-j [PK, ‘4 tN-- CH, [pK,61 CH,-k ipK.391 &H, (1) (11) m The relative acid/base strength of the resultant salts also dictates their stability to disproportiona- tion, since all salts will be acid and therefore potentially reactive towards basic formulation ad- ditives. The other essential selection issue for a salt form is the relative toxicity of the conjugate an- ion; some salts clearly fall into a desirable cate- gory, some acceptable but less desirable (both “GO”) and some undesirable (“NO GO”). A ta- ble of salts used in pharmaceutical products marketed in the U.S. up to 1974 is given in Table 1. It would seem sensible that any acid relating to normal metabolism, or present in food and drink can be regarded as a suitable candidate for prepar- ing salts. Clearly anions that cause irritancy to the TABLE 1 FDA-APPROVED COMMERCIALLY MARKETED SALTS Anion Percent a 1.26 Acetate Benzenesulfonate Benzoate Bicarbonate Bitartrate Bromide Calcium edetate Camsylate b Carbonate Chloride Citrate Dihydrochloride Edetate Edisylate ’ Estolate d 0.25 0.51 0.13 0.63 4.68 0.25 0.25 0.38 4.17 3.03 0.51 0.25 0.38 0.13 Esylate ’ 0.13 Fumarate 0.25 Gluceptate f 0.18 Gluconate 0.51 Glutamate 0.25 Anion Iodide Percent a 2.02 Isothionate ’ 0.88 Lactate 0.76 Lactobionate 0.13 Malate 0.13 Maleate 3.03 Mandelate 0.38 Mesylate 2.02 Methylbromide 0.76 Methylnitrate 0.38 Methylsulfate 0.88 Mutate 0.13 Napsylate 0.25 Nitrate 0.64 Pamoate 1 .Ol (Embonate) Pantothenate 0.25 Phosphate/ 3.16 diphosphate Polygalacturonate 0.13 Salicylate 0.88 Stearate 0.25 Glycollylarsnilate g 0.13 Subacetate 0.38 Hexylresorcinate 0.13 Succinate 0.38 Hydrabamine h 0.25 Sulfate 7.46 Hydrobromide 1.90 Tannate 0.88 Hydrochloride 42.98 Tartrate 3.54 Hydroxynaph- thoate 0.25 Teoclate ’ 0.13 Triethiodide 0.13 Cation Percent a Cation Percent a
`
`Benzathine k Chloroprocaine Choline Diethanolamine Ethylenediamine Meglumine ’ Procaine 0.66 0.33 0.33 0.98 0.66 2.29 0.66 Metallic: Aluminium Calcium Lithium Magnesium Potassium Sodium Zinc 0.66 10.49 1.64 1.31 10.82 61.97 2.95 a Percent is based on total number of anionic or cationic salts in use through 1974. b Camphorsulfonate. ’ 1,2-Ethanedisul- fonate. d Laurylsulfate. e Ethanesulfonate. f Glucoheptonate. s p-Glycollamidophenylarsonate. h N,N’-Di(dehydroabiety1) ethylenediamine. i 2Hydroxyethanesulfonate. ’ X-Chlorotheo- phyllinate. k N,N’-Dibenzylethylenediamine. ’ N-Methylgluca- mine. Reproduced from Berge et al. (1977) with permission of the copyright owner (J. Phorm. Sci.).
`
`Page 2
`
`"NO-GO"
`

`

`203 GI tract should be avoided for some types of drug, e.g. anti-inflammatories, laxative surfactant anions for anti-secretory drugs and conjugate anions with intrinsic toxicity, e.g. oxalate.
`
`Properties desired of the salt form (WANTS)
`
`mglml
`
`The desires or ‘WANTS’ of a salt form are dictated by the nature of the required dosage forms, their process and desired biological perfor- mance. Thus, it is somewhat difficult to provide a complete overall specification of ‘WANTS’ for a series of salt forms, but ideally the bulk salt should be completely chemically stable, non-hy- groscopic, not cause processing problems, and dis- solve quickly from solid dosage forms. Because of simple availability and physiological reasons, the monoprotic hydrochlorides have been by far the most frequent (- 40%) choice of the available anionic salt-forming species. Thus, there is clear precedent, and an overwhelming argument on, many grounds to immediately progress to the hydrochloride salt and evaluate other forms only if problems with the hydrochloride emerge.
`Kramer and Flynn (1972) suggest that the solu- bility of an amine hydrochloride generally sets the maximum obtainable concentration for a given amine. Many reports (Miyazaki et al., 1980, 1981) have shown that hydrochloride salt formation does not necessarily enhance the solubility of poorly solu- ble basic drugs and result in improved bioavaila- bility. This finding is based on the common ion effect of chloride on the solubility product equi- librium: BH+Cl,, KS, = BH,+, + Cl, Hydrochloride salts therefore, have the potential to exhibit a
`
`Prepare the hydrochloride; pros and cons
`
`dissolution rate in gastric fluid because of the abundance of chloride ion (0.1-0.15 M). Indeed, the Setschenow salting-out constants (k) for chloride are greatest for drugs of very low solubility (Fig. l), and can decrease the dissolution rate of the salt to below that of the free base form (Migazaki et al., 1980), which shows LOGS,
`Fig. 1. Relationship between solubility in water and salting-out constant at 25°C (left) and 37“C (right). Key: A = phenazopyridine; B = cyproheptadine; C = bromhexine; D = trihexyphenidyl; E = isoxsuprine; F = chlortetracycline; G = methacycline; H = papaverine; and I = demeclocyline. Adapted from Miyazaki et al. (1981). Reproduced with permis- sion of the copyright owner (J. Pharm. Sci.). that a precipitous drop in drug solubility occurs as the free Cl- level is increased. An example of a basic drug showing a strong chloride-ion dependence is prazosin. CH,O CH,O NH, K,,=2.2x10-6M@300C Solubility/mg.ml-t @ 30°C Hydrochloride Base 0.1 M HCl water water 0.037 1.40 0.0083 Chloride, as well as other inorganic anions have the potential to form insoluble complex salts with weak bases (Dittert at al., 1964), which are then potentially less bioavailable than the free base form. The formation of these complex salts is controlled by their stability constant K,. Drw+) = KC Drug,, + xH + + Drug . Hz (aq)
`
`reduced
`
`Page 3
`
`

`

`204 Evaluation of K, for triamterene (Tr) yields val- ues of x = 0.5 for chloride, suggesting that one proton solubilizes two molecules of the drug, i.e. the complex is Tr,H+ Cl. With hydrochloride salts there is frequently an ‘overkill’ on acid strength, which leads to a very low pH for an aqueous solution (Nudelman et al., 1974) of the salt. This can then limit the utility of hydrochloride salts in certain parenteral dosage forms, or lead to packaging incompatibilities with pharmaceutical metal containers (aerosols). Other problems frequently arise as the result of the polar nature of hydrochloride salts. Their high hydrophilic nature, favouring wettability probably as a result of the polar ionized groups being exposed on the crystal surfaces, leads to water vapour sorption (hygroscopicity) which on occa- sions, may be excessive. This can result in processing difficulties (e.g. powder flow) and re- duce the stability of a hydrolytically unstable drug. This latter effect is exacerbated by the resulting very low pH of the loosely bound moisture. These problems can be particularly acute with dihydrochlorides (or disulphates). Also, the dif- ference in the strength of the basic centres in dihydrochloride salts can lead to a gradual loss of one of the hydrochloride moieties by release of hydrogen chloride gas (Lin et al., 1972) at elevated temperatures or under reduced pressure (i.e. freeze-drying). Also, their extreme polar nature results in excessive hydroscopicity (Boatman and Johnson, 1981) eventually leading to deliques- cence. Thus, progression of a hydrochloride salt should be a first move, but if the problems with that salt form arises due to some of the reasons outlined, then the real selection issue for a salt form emerges-what trends are available for guidance?
`Each drug and its associated range of dosage forms will present different salt form require- ments, and it is perhaps best to discuss salt selec- tion further by outlining some of the trends in salt properties that may facilitate selection. The pivot of melting point A change in the development of a compound from the free base to a salt may be promoted by a need to moderate the kinetics and extent of drug absorption, or to modify drug processing. Unfor- tunately these desires may be mutually exclusive, as the balance between these properties is fre- quently pivoted around the melting point of the salt form. For example, an increase in melting point is usually accompanied by a reduction in salt solubility (the ideal solubility of a drug in all solvents decreases by an order of magnitude on an increase of 100°C in its melting point), whereas high melting crystalline salts are potentially easier to process. The increase or decrease in melting point of a series of salts is usually dependent on the control- ling effect of crystallinity from the conjugate an- ion. This is exemplified by considering an experi- mental drug candidate (UK47880) which has a basic pK, of 8, and therefore gives access to a wide variety of salt forms: CH, UK-47880 melting point. 74O C Salts prepared from planar, high melting aromatic sulphonic or hydroxycarboxylic acids yielded crystalline salts of correspondingly high melting point (see Table 2), whereas flexible aliphatic strong acids such as citric and dodecyl benzene sulphonic yielded oils. Thus, the comparative planar symmetry of the conjugate acid appears important for the maintenance of high crystal lattice forces. This is shown by the melting point of the conjugate acid being highly correlated with the melting point of the resultant salt form (Fig. 2). Therefore the highly crystalline salts are in this case best suited to reducing drug solubility. Alternatively it should also be feasible to build up crystal lattice forces of drugs with good hydro- gen bonding potential, by considering symmetry and hydrogen bonding potential of the conjugate acid. One salt series of interest is that for
`
`The pivotal issues for salt selection
`
`Page 4
`
`

`

`n
`
`C
`
`G Y E . b zoo- i
`
`205 TABLE 2 MELTING POINT OF SALTS OF EXPERIMENTAL COM- POUND (UK47880) AND THE CORRESPONDING CON- JUGATE ACID Melting point (O C) Legend Salt Conjugate Fig 2 acid UK 47880; free base pamoate (embonate) 4-hydroxynaphthalene- 1-sulphonate Salicylate 3_hydroxynaphthalene- 2carboxylate 2-hydroxynaphthalene-1 carboxylate anthraquinone-3sulphonate dodecylbenzene sulphonate mesylate citrate 74 A 235 280 G 170 190 D 156 158 C 223 220 E 145 120 B 234 225 F 20 - 113 20 20 153 epinephrine HO-CHCH2NHCH3 ‘OH OH Salt form Melting point (O C) epinephrine 157 tartrate 149 maleate 182 malate 170 fumarate 103 where small highly hydrogen bonding acids such as malonic and maleic gave higher melting salts, whereas the larger bitartrate and presumably sym- metrically unfavoured fumarate gave salts of lower melting point. Melting point and aqueous solubility Melting point and chemical stability The trends in melting point (m.p.) and aqueous solubility alluded to above are exemplified in the salts of a high melting antimalarial drug (Aghar- kar et al., 1976). The stability of organic compounds in the solid state is intimately related to the strength of the crystal lattice. Since the forces between molecules in a crystal are small compared with the energy F
`mD :: - 2 100 go- A 5 80- ‘; -lo- ,/// 60 80 100 200 300 MELTING POINT AClDl’C Fig. 2. Plot of melting point of UK47880 salts vs melting point of conjugate acids. Legend given in Table 2. -3 Salt form m.p. salt m.p. acid solubility (“0 (“C) (mg/ml) Free base 215 7.5 HCl 331 13 DL-lactate 172 17 1850 L-lactate 192 53 925 2-hydroxyethane sulphonate 251 620 Mesylate 290 300 Sulphate 270 20 The relationship between aqueous solubility (S,) and melting point is shown diagrammatically in Fig. 3, where log S, is correlated over a range of salts with the inverse of the melting point. Interestingly with this compound, the solubility of the hydrochloride salt in water is only approxi- mately twice that of the free base, whereas the low melting DL-lactate provides a 200-fold advantage over the free base in terms of solubility, which is a result in part of the reduced lattice energy.
`
`Page 5
`
`

`

`206 Fig. 3. Plot of aqueous solubility VI inverse of absolute melting point for a series of salts of a hydrophobic antimalarial drug. Data taken from Agharkar et al. (1976). required to break chemical bonds, liquefaction of the solid (and an increased frequency of molecular collisions) occurs before degradation begins. Thus the melting point of a compound can be an im- portant factor in determining stability. Degradation of solid drugs, when it is observed, usually occurs in the surface film phase and is accompanied by the formation of a liquid phase at temperatures below the normal melting point of the solid. Using this so-called ‘liquid layer’ ap- proach, Guillory and Higuchi (1962) investigated the stability of esters of vitamin A employing the following equation to determine the relationship between degradation rate and melting point. AAH[l] AH -- log K= R[T,] RT, where T,,, = normal melting point; Td = depressed storage temp. = storage temperature; K = degrad- ation rate constant; AH = heat of fusion. Thus, for a series of related compounds subject to a storage temperature Td, the logarithm of the degradation rate constant is inversely related to the absolute melting point of the compounds. Although this approach may be somewhat simplis- tic it may have utility as a method of assessing the bulk stability of non-hygroscopic salt forms. The melting point of a salt form also has some influence on its relative compatibility with drug combinations (Hirsch et al., 1978) or formulation excipients (Li Wan PO and Mroso, 1984) since it essentially controls the formation and extent of eutectic melts. As an additional aspect to the strength of crystal forces, the balance of the amorphous to crystalline nature in solid salts can dramatically affect their stability. This is exemplified by the sodium salts of ethacrynic acid (Yarwood et al., 1983). m.p. ( o C) B remaining after 9days@60°C Sodium ethacrynate Crystalline Amorphous 200 100 92 These results are consistent with the concept of an amorphous material being a highly viscous con- centrated solution and show that the stronger crystalline lattice forces result in superior solid state stability. Melting point and formulation processing Salt formation is frequently employed to raise the melting point (and crystallinity) of the drug species being processed. However, published work concerning this type of manipulation is somewhat sparse. The melting point of drug salts can dramati- cally affect their physical storage. Drugs (or salts) with low melting points generally exhibit plastic deformation (Jones, 1979) and thus during storage the stress exerted by the bulk mass on the asperity points of interparticulate contact can lead to the formation of localized welds leading to bulk ag- gregation. Also, if the sublimation temperature is low (e.g. ibuprofen, m.p. 76OC), intraparticulate voids can be bridged by sublimed drug again leading to aggregation. Thus on storage, the bulk drug salt will begin to cake and aggregate, thereby altering significantly its flow, compression and long-term dissolution properties. Melting point also has a crucial role in drug processing, in particular comminution and tablet- ing. Since low melting compounds tend to be plastic, rather than brittle, they comminute poorly, and frictional heating causes melting and deposi- tion of the drug on the screens and pins of the mill causing it to ‘blind’. For production of fine pharmaceutical powders this aspect is crucial to judging the correct level of filler to allow efficient
`
`Page 6
`
`

`

`207 manufacture using a cost-effective feed rate. Salt melting point can also have important implications for particle bonding on compression for tableting. Since bonding on compression oc- curs by point welding at the deformed or frag- mented particle surfaces, then at a fixed tempera- ture and pressure, a lower melting species would be expected to improve bonding. However, the pressure on the powder (and the eutectics formed with the other excipients) suppress the melting point further. The Skotnicky equation defines the fall in melting point (T,) with the pressure on the solid (9,) !%__L -VT dP, AH, where AHf = heat of fusion; V, = volume of solid; T = temperature, and therefore as well as those salts which are intrinsically low melting, salts of different values of AHf would be expected to have different abilities to cold weld in the compression process. If we compare, for example, the melting points and heats of fusion of the salts of an experimental drug candidate: Salt Tm (“(3 AH, (kJ,mol-t) Hydrochloride 280 56.5 Mesylate 135 20.5 Tartrate 213 63.6 Citrate 180 27.2 Phosphate 250 136.5 Acetate 180 167.9 the data suggest that the low melting point and low AH, for the mesylate salt would make it the most suitable candidate, on bonding grounds, for a direct compression tablet. Since the melting points of compounds are reduced under pressure, the solubility of salt forms would be expected to increase with increasing pressure. This can poten- tially cause the formation of solutions of the salts in the film of absorbed moisture on the surface of the drug (and excipient) particles which then may have an effect on drug bonding (Parrot, 1982) or cause the drug to adhere to the punches on com- pression (Wells and Davison, 1985). Conclusion The consideration of melting point is a key parameter in assessing the ‘viability’ of certain salt forms. In general, an increase in melting point, usually by maximizing or encouraging crystal sym- metry, leads to reduced solubility in all solvents, but generally improved stability, particularly if salt formation results in a crystalline solid, and easier formulation processing. For a specific salt form for parenteral use, i.e. where solubility and resultant pH is a major issue, a low melting point salt produced using a soluble fairly weak acid (see next section) probably made in situ is likely to be preferable. The pivot of drug solubility There are various solubility issues that can de- cide the viability of a particular salt form and it is perhaps worth addressing these separately to iden- tify trends that may aid salt selection. Aqueous solubility per se As indicated earlier, the solubility of a drug can be enhanced dramatically by salt formation (Agharkar et al., 1976). This enhancement may arise from a reduction in melting point, or from improved water-drug interactions. A good exam- ple of this is with the salts of chlorhexidine (Senior, 1973) where increased water solubility was not only produced by a lowering of melting point, but by increasing the hydroxylation of the conjugate acid. Chlorhexidine NH- salt Structure Melting Solubility point % w/v (“C) @ 20°C base above 134 0.008 dihydrochloride HCI 261 10.06 di-2hydroxy- naphthoate COzH 0.014 diacetate dilactate digluconate CH,CO,H 154 1.8 CH,CHOHCO,H _ 1.0 HO,C(CHOH),CO,H low 70
`
`Page 7
`
`

`

`B =
`+X,
`+BH;,,,
`free base of salt
`ILK,
`K,
`IL-H’
`B (\I =Bca,,
`HX,,,, = HX,,,
`H+
`
`I~~t; move
`
`208 The above data exemplify the importance of considering the hydrophilic nature of the con- jugate anion, as well as its role in controlling crystallinity, when considering the potential solu- bility of salts. Reduced aqueous solubility may occasionally be a crucial development factor for a drug, e.g. for an organoleptically acceptable or chemically sta- ble suspension. Such systems demand salts of low solubility, but recent experience with a series of purposely designed insoluble salts of an experi- mental drug candidate also highlighted the need to consider the solubility and pK, of the conjugate anion. RH’ x (s)
`
`ionic equmnria snow that even sparing solubility of the salt means that the level of the conjugate anion in solution will depend markedly on the pH of the fluid. Consideration of the above for a pamoate salt, which has pK,‘s of the parent acid of 2.5 and 3.1 and virtually insoluble un- ionized form, indicates that solutions of pH 5-6 will drive the equilibria to the right, with full precipitation of the free acid HX,,, and liberation of a full component in solution of the ionized base (BH’). However, if we consider the hydroxynaph- thalene sulphonic acid (PK. = 0.11) then this sys- tem provides ‘insolubility’ over a much wider pH range and is therefore far more tolerant to fluctua- tions in the fluid pH. The above aspect is important when consider- ing the potential use of ‘insoluble’ salts (e.g. pamoate) to control the absorption of a drug candidate. For example, the in vitro dissolution rates of the dimaleate and pamoate salts of a drug candidate were compared in simulated gastric and intestinal fluid. The dissolution rates were essen- tially identical in the former fluid, with rapid deposition of the pamoic acid and liberation of the free base, whereas in the latter the pamoate salt exhibited a much slower dissolution rate than the maleate. Therefore ‘control’ on the drug ab- sorption (and toxicity) may then depend on the duration of gastric residence and the pH of the gastric contents. Thus aspects such as food vs the fasted state are also important. In fact in this case, the bioavailability in the dog of the two salt forms when dosed orally from a standard capsule formu- lation were of the same order; 24% for the pamoate and 39% for the maleate. Usually it is the dissolution rate of a drug which is of major importance to the formulation and as a rule a salt exhibits a higher dissolution rate than the base at an equal pH, even though they have the same equilibrium solubility. This latter effect, which is exemplified by theophylline salts (Nelson, 1957) is due to the salt effectively acting as its own buffer to alter the pH of the diffusion boundary layer, thereby increasing the apparent solubility of the parent drug in that layer. Thus, administration of basic drugs as their salt forms (e.g. tetracycline hydrochloride) ensures that stomach emptying rather than in vivo dissolu- tion will be the rate-limiting factor in their absorp- tion. It is also possible that increased drug absorp- tion may occur with salts due to their effect on the surface tension of the gastrointestinal fluids (Berge et al., 1977). Salt solubility and pH of salt solutions Enhancement of the aqueous solubility of a drug by salt formation can occur due to dif- ferences in the pH of the saturated salt solutions. A soluble acid salt of a weakly basic drug will cause the pH to drop as the salt is added to the solution. This pH drop will, in turn cause more drug to dissolve, and this process will continue until the pH of maximum solubility is reached (see Fig. 4). The equilibrium solubility(ies) are then given by: s =
`
`s+(l
`
`+
`
`lopn-px~) for pH = pH,,, i.e. when the ionized form is solubility limiting and s = S,(l + 10pk*PP”) for pH = pH,,, where the unionized form is solubility limiting and pH,, is given by the solu- tion of the equality of pH for the above two
`
`Page 8
`
`experimental drug
`

`

`4
`
`5
`
`6
`
`SC
`
`pH~x\p~P;Si(l
`
`8
`
`9
`
`7
`P”
`
`;IOpKamPH)
`
`10
`
`11
`
`209 1 , 1
`Fig. 4. Solubility of A in water at ambient temperature ( - 23°C) as a function of pH. All data arc in mg/ml calculated in terms of free base equivalent. The lines drawn through the data are theoretical and were calculated using 0.067 mg/ml as the free base solubility, 11.5 mg/ml as the hydrochloride solubility and 8.85 as the pK,. Data by both gravimetric (m) and GLC (0) procedures were in good agreement. Adapted from Kramer and Flynn (1972) with permission of the copyright owner (J. Phurm. Sci.). equations where PH,,
`
`=pK,+log y and implies that both free base and salt form can exist simultaneously in equilibrium with the saturated solution. Thus, large pH shifts on dissolution of salts suggests that a large amount of conjugate acid is dissociating and therefore, a relatively high solu- bility is then obtained. If we consider physiologi- cal pH, a low pK, for a conjugate acid of high aqueous solubility, would appear to give the best change of obtaining the lowest pH,, and the highest aqueous salt solubility. For example, the solubilities of a series of salts of a drug candidate and the pH of the saturated solutions were as follows: Type Salt PH,, Conjugate acid Solu- pK, Solu- bility bility (mg/ (mg/ ml) ml) Hydro- chloride Mesylate Tartrate Citrate Phos- phate Acetate 2.71 35.9 -6.1 2.57 51.2 ~ 1.2 4.21 0.49 3.03 1470 3.30 2.16 3.13 2400 5.31 10.31 2.15 5.29 8.04 4.76 indicating that the salts of stronger acids (HCl, methane sulphonic) produce the lowest slurry pH and the highest salt solubility. In this case the solubility and resultant low pH of the hydrochlo- ride is suppressed by the common ion effect. The solubility of salts such as the lactate (PK. of conjugate acid is 3.85, with infinite solubility) may offer significant advantage over for example the acetate, tartrate and citrate. The pH of a salt solution can be a deciding factor in the selection of a salt for a parenteral dosage form. Ideally to avoid pain on injection the pH of i.v. parenterals should be between pH 3 and 9, and so highly acidic salts such as the hydrochlo- ride and mesylate are probably best replaced by an acetate salt. Kramer and Flynn (1972) have shown that for a series of hydrochloride salts, by making analysis of the differential heat of solutions of the ionized and unionized species, that the temperature de- pendencies of the solubilities of the hydrochloride salts were considerably lower than those of the corresponding free base form. This may have im- portant implications for solution dosage form des- ign and storage conditions. Salt solubility and salt stability As well as the relationships between salt melt- ing point and stability raised earlier, it is also clear that low solubility and low hygroscopicity can contribute significantly to the stability of a salt form. The former aspect is obviously important in developing a stable aqueous suspension formula- tion of a hydrolytically unstable water soluble
`
`Page 9
`
`

`

`The pivot of salt hydrophobic@
`
`210 drug; e.g. penicillin-benzithine. For salts of weak bases, the moisture associated with the bulk can be very acidic (the salt will buffer the available moisture), and can potentially cause severe hydrolytic degradation of the parent drug. Classic examples of this phenomenon are thiamine salts (Yamamoto et al., 1956, 1957) where the stability is related to their hygroscopicity, aqueous solubility and the resulting pH. Thus, to improve drug stability by salt forma- tion, it is clearly not only important to control hygroscopicity, but also to consider carefully the strength of the conjugate acid used to form the salt. This is particularly important for compacted dosage forms where salt and excipient share the available moisture, particularly when the majority of the available moisture comes from the excipient rather than the drug. Thus, assessment of salt stability in compressed and non-compressed sys- tems is an important activity in preformulation studies. However, in selection terms salts of mineral acids will produce a lower pH, and higher solubility in the available moisture and therefore produce a more hostile stability environment than that from a sulphonate or carboxylate type salt. It is also apparent that another consideration in the relationship of salt stability is the hydrophobic portion of the conjugate acid. This is exemplified with xilobam (Walking et al., 1983) where aryl sulphonic acids salts were prepared to protect this easily hydrolyzed base. (-JNHCONQ 6 pJSO;” - CH, S 04 xilobam tosylate 1.2-napsylate 168OC 177OC. 177OC saccharinate 150°C The rationale behind the choice of these salt forms was that they comprise fully ionized acids and therefore present pH-independent aqueous solu- bility in biological fluids. However, as opposed to the poorly stable hydrochloride and sulphate salts, the aryl groups present a hydrophobic barrier to minimize hygroscopicity and dissolution in the surface moisture. The highest melting (least solu- ble?) salt, l-napsylate, proved to be the most stable, with full retention of potency following 7 days storage at 7O”C/74% RH, whereas only 18% of the base remained after this challenge. The I-napsylate salt also provided full and rapid dis- solution in vitro (t,,,, 15 mm).
`Although the xilobam example above serves to demonstrate that one has to consider the hy- drophobicity of the conjugate anion to control salt stability, it is clear that this property is pivotal on two others; salt hygroscopicity and wettability. Thus, once again, a balance of salt properties is required so that hygroscopicity is not reduced at the gross expense of salt wettability leading ulti- mately to dissolution rate and bioavailability problems. Hydrophobicity und hygroscopicity Soluble ‘polar’ salts have