Isothermal Colorimetry

2005, 6, 215-222 215

1389-2010/05 $50.00+.00 2005 Bentham Science Publishers Ltd.

Application of Solution Calorimetry in Pharmaceutical andBiopharmaceutical Research

P.G. Royall1,* and S. Gaisford2

1Department of Pharmacy, Kings College London, Franklin-Wilkins Building, 150 Stamford Street, London, SE1 9NHand 2Department of Pharmaceutics, School of Pharmacy, University of London, 29/39 Brunswick Square, London,WC1N 1AX, UK

Abstract: In solution calorimetry the heat of solution (DsolH) is recorded as a solute (usually a solid) dissolves in anexcess of solvent. Such measurements are valuable during all the phases of pharmaceutical formulation and the number ofapplications of the technique is growing. For instance, solution calorimetry is extremely useful during preformulation forthe detection and quantification of polymorphs, degrees of crystallinity and percent amorphous content; knowledge of allof these parameters is essential in order to exert control over the manufacture and subsequent performance of a solidpharmaceutical. Careful experimental design and data interpretation also allows the measurement of the enthalpy oftransfer (DtransH) of a solute between two phases. Because solution calorimetry does not require optically transparentsolutions, and can be used to study cloudy or turbid solutions or suspensions directly, measurement of DtransH affords theopportunity to study the partitioning of drugs into, and across, biological membranes. It also allows the in-situ study ofcellular systems. Furthermore, novel experimental methodologies have led to the increasing use of solution calorimetry tostudy a wider range of phenomena, such as the precipitation of drugs from supersaturated solutions or the formation ofliposomes from phospholipid films. It is the purpose of this review to discuss some of these applications, in the context ofpharmaceutical formulation and preformulation, and highlight some of the potential future areas where solutioncalorimetry might find applications.

Key Words: Solution calorimetry, polymorphism, amorphous content, pharmaceuticals, dissolution, enthalpy of transfer.

1. INTRODUCTION

Many active pharmaceutical ingredients (APIs) exist inthe solid-state or are formulated into solid dosage forms.This confers several advantages; stability in the solid-state isusually greater than in the liquid state, solids are easier totransport, package and store and the majority of medicines,at least in the United Kingdom, are formulated as tablets asthe oral route is both quick and convenient. However, solid-state formulations are usually slower-acting than liquidmedicines containing the same active, because the first eventthat must occur following administration is that the API mustdissolve in a suitable biological fluid (such as saliva orgastric juices). On a molecular level, this requires theintramolecular forces holding the solid lattice together to beovercome (endothermic) and the formation of new inter-actions with the solvent (generally exothermic). This event is(usually) rate-limiting, and thus controls the observeddissolution rate of the API from the formulation.

The intra-molecular forces in a crystalline material (thelattice energy) will vary if the material exhibits polymor-phism, (i.e. if the molecules can pack in more than onecrystalline formation), the most stable polymorph having thehighest lattice energy. It is therefore imperative that thecrystal form of any solid-state API is known, since thedissolution rates of each polymorph will be different.

*Address correspondence to this author at the Department of Pharmacy,Kings College London, Franklin-Wilkins Building, 150 Stamford Street,London, SE1 9NH, UK; Tel: +44 (0)20 7848 4780; Fax: +44 (0)20 78484800; E-mail: [email protected]

Moreover, over time all the metastable polymorphs willconvert to the stable polymorph, which may have adisastrous effect on the efficacy of the formulated product. Aclassic example of this is provided by chloramphenicolpalmitate, which has three polymorphs, termed A, B and C[1]. The A polymorph is unstable and is not allowed to beincluded in any medicine; similarly, the C polymorph hasvirtually no appreciable bioavailability and its percentage inthe final product must be strictly controlled. Any formulatorof chloramphenicol palmitate must therefore know, and beable to quantify, the amount of the B polymorph in anyformulation.

In addition to polymorphism, solid APIs may also exhibitamorphicity (a lack of long-range molecular order). Becauseamorphous materials dont have a lattice energy and areessentially unstable (over time they will recrystallise to acrystalline form) they usually have appreciably fasterdissolution rates than their crystalline equivalents, whichmakes them especially suited for formulation into fast-actingmedicines. The accidental inclusion of amorphous contentin what is otherwise presupposed to be a crystalline materialis also a hazard associated with processing of solid-state

It is often stated that amorphous materials have higher solubilities than theircrystalline counterparts; since solubility is an equilibrium state, this cannot be true.What may often happen is that the fast dissolution rate of the amorphous form resultsin the production of a super-saturated solution which over time, although notnecessarily in the time frame involved in drug administration and absorption, willprecipitate to form a saturated solution. As such, it should be stated that amorphousmaterials may exhibit higher apparent solubilities than their crystalline equivalents.

216 Current Pharmaceutical Biotechnology, 2005, Vol. 6, No. 3 Royall and Gaisford

materials (during, say, milling or compression steps) [2-3].Although the percentage of amorphous content introduced inthis way is usually low (of the order of 1% w/w) its locationprimarily on the surface of what are usually small particlesgives it a disproportionate control on the properties of thematerial [4]. It is clear, then, that the detection and quanti-fication of (often small) amorphous contents is of the utmostimportance during the characterization of a solid pharma-ceutical.

One technique that offers the sensitivity to differentiatebetween polymorphs as well as to detect low (typicallyaround 1% w/w) amorphous contents is solution calorimetry.The principle of solution calorimetry is simple; the heatchange when a small quantity of a solid or liquid sample isdispersed into a (relatively) large volume of solvent ismeasured (either directly or by measuring a temperaturechange which is subsequently converted to a heat change).For a pure material dissolving into a solvent this results inthe enthalpy of solution (D solH), an enthalpy change thatreflects contributions from the bonds broken when thecrystal lattice breaks apart (D latticeH) and from the bondsformed when the molecules are solvated (D solvationH). Thiscan be expressed as;

HHH solvationlatticesol D+D=D Equation 1

Depending on the relative magnitudes of D latticeH andD solvationH, the heat of solution can have a positive (endo-thermic) or negative (exothermic) sign. Careful experimentaldesign (discussed in more detail below) is required to ensurethat the measured enthalpy does not contain erroneouscontributions from other sources, such as dilution (liquidsamples), friction effects (ampoule breaking and samplestirring) and sample-sample interactions (if an ideal solutionis not formed), or simply as a corollary of poor experimentaldesign or execution. Since D solH differs for each polymorphof a solid sample, it is clear that solution calorimetryprovides valuable information that allows the detection andquantification of polymorph content.

Moreover, since D solH contains information on the attrac-tive forces holding a solid together, solution calorimetry canprovide fundamental information on the nature of multi-component systems. Such systems include solid-dispersions,polymeric systems and proteins freeze-dried with carbo-hydrates. As such, solution calorimetry is a powerful toolthat can be used to investigate a wide range of pharma-ceutical systems, and it is the purpose of this review todiscuss some of these applications in the context of thepreparation and formulation of novel pharmaceuticals andbiopharmaceuticals.

2. THE PRINCIPLES OF SOLUTION CALORIMETRY

A solution calorimeter records the heat change when asample (solid or liquid) is dissolved in a large volume ofsolvent. There are two types of solution calorimeter designcommercially available; instruments that operate on a semi-adiabatic principle (i.e. that record a temperature changeupon reaction) and instruments that operate on a heat-conduction principle (i.e. that record a power change directlyupon reaction). The sensitivities of these instruments, andhence the quantities of solute and solvent required for

experiment, vary considerably (typically, semi-adiabaticinstruments are less sensitive and require much larger samplevolumes).

In a carefully constructed experiment, the enthalpy ofsolution, D solH, is measured directly. Care must be taken toensure that other contributing heat sources are accounted for.These may include effects from breaking the ampoule andsample stirring (most easily compensated for by subtractinga blank experiment from the sample experiment), changes inthe solvent activity because of solute dissolution, changes inthe rate of evaporation of solvent into the headspace andchanges in volume upon mixing of the two phases. The lastthree effects can be assumed to be negligible if an ideal(dilute) solution is formed; if this is not the case, a fact thatwould be confirmed by obtaining different values of D solHwith different amounts of sample, then the data must beextrapolated back to infinite dilution to give D solH

.

2.1. Types of Solution Calorimeter

Semi-adiabatic and heat-conduction instruments operateon different principles and, while these have been discussedin detail elsewhere (See for example, 5-7) it is worth brieflyconsidering these differences here because they impact uponthe subsequent discussion of the data.

2.1.1. Semi-Adiabatic Solution Calorimeters

In an ideal adiabatic calorimeter there is no heat ex-change between the calorimetric vessel and its surroundings.This is usually attained by placing an adiabatic shield aroundthe vessel. Thus any change in the heat content of a sampleas it reacts causes either a temperature rise (exothermicprocesses) or fall (endothermic processes) in the vessel. Thechange in heat is then equal to the product of the temperaturechange and an experimentally determined proportionalityconstant (or calibration constant, e ). The proportionalityconstant is usually determined by electrical calibration.Thus;

eq

T =D Equation 2

eF=

t

T

d

d Equation 3

where F represents power. Ideally, the value of e returnedafter calibration should equal the heat capacity of thecalorimeter vessel (Cv, the vessel including the calorimetricampoule, block, heaters, thermopiles and the sample) but inpractice losses in heat mean the value may differ slightly.However, assuming the losses are the same for both sampleand reference, the power value returned will be accurate. It isalso the case that Cv varies depending on the heat capacity ofthe sample being studied, which affects the measuringsensitivity of the instrument. Thus, a small heat capacityresults in a large rise in temperature for a given quantity ofheat and, consequently, better sensitivity. However, calori-meters with low heat capacities are more sensitive toenvironmental temperature fluctuations and have lowerbaseline stabilities. Any calorimeter design therefore resultsin a compromise between baseline stability and measurementsensitivity.

Application of Solution Calorimetry in Pharmaceutical Current Pharmaceutical Biotechnology, 2005, Vol. 6, No. 3 217

In practice, true adiabatic conditions are difficult toachieve and there is usually some heat-leak to the surround-ings. If this heat-leak is designed into the calorimeter (suchas is the case with the SolCal, an example of a commercialsemi-adiabatic calorimeter, Thermometric AB, Jrflla,Sweden) the system operates under semi-adiabatic (orisoperibol) conditions and corrections must be made in orderto return accurate data. These corrections are usually basedon Newtons law of cooling (the most common being themethod of Regnault-Pfaundler, discussed below).

In the case of the SolCal (the principles apply to allsimilar solution calorimeters), at the start of an experimentthe instrument is held above or below the temperature of itsthermostatting bath (typically by up to 200 mK). With timethe instrument will approach the temperature of thethermostatting bath; data capture is initiated when thisapproach becomes exponential (this assumption is anecessary precursor to employing the heat-balance equationsused to calculate the heat evolved or absorbed by the systemcontained within the vessel). Thus, the response due todissolution, and any electrical calibrations (usually two areperformed; one before and one after the break to ensure theheat capacity of the system has remained constant), must beperformed before the instrument reaches thermal equilibriumwith the bath. In practice, this limits the technique tostudying events that, ideally, reach completion in less than30 min.

Upon completion of an event in a solution calorimeter aquantity of heat will have been recorded. As noted above, theheat will be given by the product of the temperature changeand the calibration constant (which in the ideal case is theheat capacity of the vessel). This interpretation assumes thatthe measured temperature change arises solely from theevent occurring in the vessel; in practice, other events, suchas ampoule breaking, stirring and heat-leakage, all contributeto the temperature change of the vessel. For accurate dataanalysis these effects must be removed from the observedtemperature change ( D Tobs) to give D Tcorr, the temperaturechange that would have occurred under ideal conditions.Thus;

adjcorrobs TTT D+D=D Equation 4

where D Tadj is defined as the temperature change arisingfrom all the other contributing events in the vessel. Usuallythe method of Regnault-Pfaundler, which is based on thedynamics of the break, is used to determine the value of D Tadj[8]. In this case;

D Tadj =1t

(T - T)dttstart

tend

Equation 5

where T is the temperature of the vessel and its contents attime t, T is the temperature that the vessel would attain afteran infinitely long time period, tstart and tend are the start andend times of the experiment respectively and t is the timeconstant of the instrument. Note that T is effectively thevalue of T at t and is commonly described as the steady-state temperature of the vessel. The time constant has unitsof seconds and can also be expressed as;

k

et = Equation 6

where k is the heat exchange coefficient of the vessel.

The values of T and are calculated by analysis of thebaseline regions immediately preceding and following thebreak. These baseline sections will be approaching thetemperature of the surrounding heat-sink exponentially andare described by;

T = T + (T0 - T )e

- t

t Equation 7

The data are then fitted to Equation 7 using a least-squares minimising routine to return values for T and t .Once these are known, D Tadj can be calculated. This value isthen used to calculate D Tcorr for the sample break and also forthe two electrical calibrations (D Tcorr, calibration). There shouldbe no significant difference in the D Tcorr, calibration valuesdetermined for the two calibrations and an averaged value isused. The calibration constant is then determined from;

ncalibratiocorr,

ncalibratio

T

Q

D=e Equation 8

the heat-change for the break is then easily determined;

corrreaction . TQ D= e Equation 9

2.1.2. Heat-Conduction Calorimeters

A heat-conduction calorimeter is surrounded by a heat-sink, which acts to maintain the system at a constant tem-perature. Between the vessel and the heat-sink is a thermo-pile wall. Any heat released or absorbed upon reaction isquantitatively exchanged with the heat-sink. The thermopilesgenerate a voltage signal that is proportional to the powerflowing across them; this signal is amplified, multiplied bythe cell constant (determined through electrical calibration)and recorded as power versus time. An isothermal system isnot limited to reaction processes that reach completionwithin 30 min, as semi-adiabatic instruments are, because itis always (essentially) in equilibrium with its surroundingheat-sink. Furthermore, the greater measuring sensitivity ofthe thermopiles (as opposed to the thermisters used in semi-adiabatic instruments) means that smaller sample masses canbe used.

A further consideration of the use of isothermalinstruments concerns dynamic correction of the data. Theaim of dynamic correction is to remove the effect of thethermal inertia inherent in any calorimeter (i.e. the delaybetween heat being released by the sample and that heatcausing a measurable voltage to be generated by thethermopiles) and it is principally used for short-term events,typically in titration experiments. However, since oneoutcome from dynamic correction is an improvement in thesignal to noise (S/N) ratio of the data it offers the potential toreduce the standard deviation of dissolution experiments,because peak areas can be determined with greater precision.

In the case of a typical instrument (the 20 ml microreaction ampoule, Thermometric AB, Sweden), dynamiccorrection is achieved by application of a modified form ofthe Tian equation;


2R

2

21R

21RC d

d..

d

d).(

t

P

t

PPP tttt +++= Equation 10

where PR and PC are the raw and corrected powersrespectively and t 1 and t 2 are termed the first and secondtime constants of the instrument (for a further discussion ofthe derivation and use of Equation 1 see, for example, [9]).The time constants, t 1 and t 2, are determined by a leastsquares analysis of data following an electrical calibration(and are hence not user defined). It is important to note that anumber of assumptions are made in order to derive Equation10 (the major assumption being that there are no temperaturegradients within the sample) and that it only approximatesthe true dynamic delay inherent to the instrument. Thecorrected data so produced, while much more closelyresembling the true response of the sample, therefore oftencontain artefacts, such as overshoots where both endo- andexothermic events are indicated even though it is known thatonly one event is occurring in the sample. In principle, theseartefacts could be removed by altering the values of t 1 and t 2but this is difficult in practice. It is therefore easier to usecorrected data to determine reaction enthalpies only and tonote that the use of such data to elucidate kinetic informationmust be undertaken with caution. It can be shown that thetotal net heat change recorded for both dynamicallycorrected and raw data, in the ideal case, is the same [9].

2.1.3. Sample Solvent Mixing

There are several methods by which the solute andsolvent can be introduced. The solute and solvent can beheld in separate chambers of the same cell - rotating thesample cell end-over-end mixes the materials and initiatesthe interaction (this is the system employed by, for instance,the C80, Setaram, France); the solute can be held in a sealedglass ampoule which is broken into the solvent to initiatereaction (used in the 2225 precision solution calorimeter,Thermometric AB, Sweden); the solute can be held inreusable metal canisters which are broken into the solvent(used in the 20 ml micro solution ampoule, ThermometricAB, Sweden); or liquid solutes/solutions can be titrated intoa cell containing solvent (such as the VP-ITC, Microcal Inc.,USA).

2.2. Calibration

Calibration is vital to ensure that instruments are opera-ted properly and are functioning correctly, that data fromdifferent instruments or different laboratories are comparableand, perhaps most importantly when considering pharma-ceutical samples, that the data obtained are validated and canbe incorporated into regulatory documents. Most calori-meters are calibrated using an electrical substitution method.In this case, a resistance heater (usually located under thesample cell) produces a known amount of heat, which ismeasured by the instrument. The heat output recorded isadjusted until the measured and expected heats are the same(by multiplying the raw data signal by a constant value; thecell constant). However, it is debatable whether an electricalcalibration truly represents an accurate method by which tovalidate instrument performance, primarily because thesource and rate of heat generation will be different from thatcaused by a chemical interaction. In this sense, chemical

standards offer a better alternative and have been the subjectof some discussion in the literature.

There are a number of requirements imposed upon a testreaction; it should be robust, simple to perform, requirecommonly available materials that require no specialpreparation prior to use and it should be applicable across arange of instrumentation [10]. A number of chemical testreactions for solutions calorimeters have been proposed anddiscussed, including the dissolution of Tris in 0.1 M HCl [5,11], the dissolution of KCl or NaCl in water [12-14] and thedissolution of propan-1-ol in water [15]. The dissolution ofsucrose in water can also be used [16-17], although this isnot currently recognised as a test reaction.

Of these systems, the dissolution of KCl into water isusually recommended because it is robust, easy to performand a standard reference material is available from NIST (theNational Institute for Standards and Technology, USA); theenthalpy of solution of the NIST certified KCl into water is17.584 0.017 kJ mol-1 [13]. However, the use of KCl is notwithout drawbacks; principally, the value of D solH varies as afunction of the concentration achieved after dissolution,because of the effect of the enthalpy of dilution (F L); thus,the certified value for the NIST reference material of 17.584kJ mol-1 applies only if a final concentration of 0.111 mol kg-1 is attained in the calorimetric vessel. This corresponds to amolar ratio of water to KCl of 500 to 1 and is often writtenas D solH (500 H2O, 298.15 K). If measurements are per-formed under different conditions, then the value obtained(nH2O, 298.15 K) must be corrected to that which wouldhave been recorded at 500 H2O, in order to draw com-parison. These corrections are explained in the certificationcertificate supplied with the NIST sample [13], although thedata supplied there apply only to experiments performedwhere n varies from 100 to 1000.

The effects of KCl concentration on D solH have beenstudied extensively by Kilday [12], who corrected theobserved enthalpy values over a range of water ratios (n =500 to 10000) to account for the enthalpy of dilution; thisresulted in a value for D solH , the enthalpy of solution atinfinite dilution (D solH = 17.241 0.018 kJ mol

-1).

Modern solution calorimeters often use microgramsamples and are capable of detecting very small powers; oneconsequence of this is that it is not possible to perform theKCl experiment under the NIST certification conditionsbecause the heat generated would be of a magnitudesufficient to saturate the amplifiers. Because of this severalrecent studies have been conducted looking at the applicabi-lity of test reactions to modern solution calorimeters [7, 18].The outcome from these studies appears to be that if KCl isto be used then it is a smaller, and easier, correction to theenthalpy of solution at infinite dilution, but that sucrose mayoffer a better, and cheaper, alternative, especially for heat-conduction instruments that use very small (mg) samples.

3. APPLICATIONS

3.1. Polymorphism

As mentioned above, if a drug is prepared in the solid-state then it is highly likely it will exhibit polymorphism andit is imperative that the polymorph(s) present in a sample are


known before formulation, such that the resulting medicineperforms within its pharmacopoeial specifications. For newdrug entities, the pressure to bring a product to marketusually means that the developing company doesnt have thetime to characterise the relative bioavailability of eachpolymorphic form; usually, the most stable form is selectedfor development and the formulation is tailored to ensure itsstability. This formulation strategy ensures compliance withthe guidance provided by the International Commission onHarmonisation (ICH) on the selection of solid forms of drugs[19]. Similarly, for Abbreviated New Drug Applications(ANDA) the sponsor must provide evidence that the pro-posed generic product and the original product are pharma-ceutically equivalent and bioequivalent [20]; clearly, nothaving any control or knowledge of the polymorphic formspresent would render submission of such data impossible.

Solution calorimetry allows the direct measurement ofthe lattice energy of a sample, which alters with each poly-morphic form of the material and, although other analyticaltechniques may be used to discriminate between polymorphs(such as DSC, XRPD, solid-state NMR, IR and Ramanspectroscopy and microscopy [21]), the technique has threeprinciple advantages; firstly, the measurement is performeddirectly at the storage (or, indeed, any desired) temperature,in contrast to DSC measurements which usually need to beextrapolated to give an indication of structure at roomtemperature; secondly, the data allow comparison betweensamples, giving insights into batch-to-batch or formulationvariability; and thirdly the various contributions to themeasured heat of solution can be dissected out by carefulexperimentation, allowing a mechanistic analysis of theprocess under investigation. The principal drawback, relatedto the latter point above, is that the heat of solution is acomposite of all those processes that are occurring duringdissolution and it may sometimes be impossible to separatethe heat into its component parts.

Solution calorimetry has been used to assess the differentpolymorphs of many drugs. For instance, forms I and II ofcyclopenthiazide have been shown to have comparableenthalpies of solution (~6 kJ mol-1) while form III has ahigher enthalpy of solution of 15 kJ mol-1 [22]. Similarly,solution calorimetry has been used to characterise the threepolymorphic forms of a pre-clinical drug, Abbott-79175 [23]and the polymorphs of an angiotensin II antagonist agent(MK996) [24].

A study of urapidil resulted in the enthalpies of solutionof forms I and II of the drug (21.96 and 23.89 kJ mol-1

respectively) being determined [25]. However, the authorsnote that urapidil exists in an additional form (form III),which they were unable to obtain in a pure form. By mea-suring the heat of solution of the mixture (D mixH), and usingDSC data to determine the percentage of each polymorph ina sample containing form III (7.4% form I, 2.7% form II and89.9% form III), they calculated the enthalpy of solution ofform III using the following relationship;

).().().( IIIIIIIIIIIImix HXHXHXH D+D+D=D Equation 11

where XI, XII and XIII are the fractions of form I, II and IIIrespectively and D HI, D HII and D HIII are the heats of solution

of forms I, II and III respectively. It was determined thatD HIII equaled 22.98 kJ mol

-1.

Since Hesss law allows the interpretation of D solH byinvoking any number of steps, as long as the starting state isthe solid and the final state is the solution, one approach tothe analysis of dissolution data is to consider the heat ascomprising two major contributions; one which representsthe breaking of the solid-state bonds and the other whichrepresents all the processes through which the moleculesbecome solvated [26]. If bond breaking is considered to besimilar to a melting process, then it can be represented byD fH

298, the enthalpy of formation of a supercooled liquid at298 K (assuming the dissolution experiment is conducted at298 K), while the enthalpy changes associated with mixingand solvation can be represented by D m H. This can berepresented as shown in Fig. 1.

Fig. (1). A description of dissolution, based on Hesss law [26].

This model has been applied to the study of the dissolu-tion of various poly(ethylene) glycols, commonly encoun-tered pharmaceutical excipients, in water [26], where D fH

298

was measured using DSC data and D solH was measured inthe solution calorimeter. Four different molecular weightPEGs were selected (3, 400, 6000, 10, 000 and 20, 000) andsamples were prepared with different thermal histories;samples were quench-cooled from the melt, slow-cooledfrom the melt (5 oC min-1) or left untreated. It was found thatD solH varied as a function of both molecular weight andthermal history, although no discernable relationships couldbe found. According to the model above, D mH should beconstant for each molecular weight PEG, independent of anythermal history, since this represents the dissolution ofindividual molecules, while D fH

298 should be highly depen-dent on thermal history. Accordingly, the D soHl values willvary, since this parameter is the sum of D fH

298 and D mH.Subtraction of the D fH

298 values, determined from DSC data,from the measured D solH values showed that, within a 5%confidence level the D mH values did remain constant foreach molecular weight PEG, and that the measuredvariabilities in the measured heats of solution derived fromchanges in D fH

298.

The correlation of the heats of solution of differentpolymorphs of a drug with dissolution rates has also beenattempted. For instance, Terada et al. [27] showed there wasa linear correlation between the heats of solution of differentpolymorphs of indomethacin and the logarithms of their


initial dissolution rates (determined by the rotating diskmethod). The same authors showed a similar relationship fora range of samples of terfenadine of varying crystallinity.

In some cases, a drug will be insoluble in water, and inorder to derive some meaningful information on its poly-morphic forms it is desirable to measure heats of solution intwo non-aqueous solvents. Again, in accordance with themodel discussed above, the value of D solH for each poly-morph will alter in each solvent. However, the differencebetween the heats of solution for each polymorph should beconstant irrespective of the solvent used;

ConstantTII form sol,I form sol, =D=D-D HHH Equation 12

where D TH is the transfer enthalpy, and is the energyrequired to transfer between the two forms. Differences inD solH determined in each solvent will reflect changes in thesolvation interaction with that solvent.

This approach has been used to investigate the poly-morphs of enalapril maleate, by recording the heats ofsolution of forms I and II of the drug in acetone and metha-nol [28]. The differences in the heats of solution of forms Iand II in each solvent (2.88 and 2.13 kJ mol-1 in acetone andmethanol respectively) were the same within experimentalerror, and gave D TH, the energy that would be required forform II to convert to form I. The difference in the heat ofsolution for forms I and II in methanol, compared withacetone, was 23.53 kJ mol-1; that is, the heat of solution inmethanol was more exothermic than in acetone. This valuewas found to be consistent with the formation of a singlehydrogen bond.

Similar data, recorded using ethanol and methanol, havebeen recorded for terfenadine [29], where a difference inenergy of approximately 13 kJ mol-1 was found betweenforms I and II. In this case, it was observed that D solH for anysolid form of terfenadine was around 4 kJ mol-1 lower inmethanol than ethanol, indicating that polar groups on thedrug molecule play an important role in the solute/solventinteraction.

A different approach to study water insoluble drugs is touse a concentrated surfactant solution instead of pure wateras the solvent. For instance, the enthalpies of solution of anumber of cimetidine polymorphs have been studied usingconcentrated SDS and Tween 20 solutions [30].

3.2. Determination of Degree of Crystallinity/AmorphousContent

Solid pharmaceuticals may often not be entirelycrystalline and contain regions of amorphous content. Anamorphous material (most simply described as a materialwith the structure of a liquid and the viscosity of a solid) ischaracterised by a lack of any long-range crystal structureand, because it therefore has no lattice energy, the lack of amelting point (upon heating, an amorphous solid willrecrystallise and it is the crystalline form that subsequentlymelts). Amorphous solids are thermodynamically unstableand usually exhibit fast dissolution rates compared with theircrystalline equivalents. Pharmaceutically, this is a big advan-tage for medicines that are required to be fast acting and

many solid drugs are prepared in the amorphous statealthough, as noted earlier, knowledge of the percent amor-phous content and its stability are vital prerequisites toensure the stability and continued efficacy of the medicineupon storage.

The use of solution calorimetry to quantify the degree ofcrystallinity in a solid sample is predicated on therelationship shown in Equation 11, where the measured heatof solution is given by the sum of the enthalpies and weightfractions for the crystalline and amorphous states present.The usual methodology is to prepare a calibration line ofD solH against degree of crystallinity using a number ofknown standards (usually prepared by blending theappropriate mass quantities of wholly amorphous and whollycrystalline material); the calibration plot should be linear andcan thus be used to determine the degree of crystallinity inan unknown sample.

There are two potential drawbacks to using this approach.Firstly, the sample may exhibit polymorphism, in which casethere may be more than one crystalline form present; if thisis the case then, for the calibration plot to be linear, it mustbe ensured that the crystalline material used to prepare thestandards and the unknown sample contain the sameproportions of the polymorphs. Secondly, it is likely that theinteractions in a particle that has a crystalline core andamorphous material on its surface (the likely situation for aprocessed pharmaceutical) differ from those of whollyamorphous and wholly crystalline particles, which mayresult in the calibration plot producing spurious results.

The first use of solution calorimetry for the quantitativemeasurement of the degree of crystallinity of pharma-ceuticals was by Pikal et al. [31], who measured the heat ofsolution of various b -lactam antibiotics. Subsequent studiesinclude the analysis of sulphamethoxazole from differentsources [32], sucrose [16] and clathrate warfarin sodium[16].

In many of these studies, the aim was to quantify relat-ively large mass percentages of crystalline material, givingheat changes easily within the detection limit of the tech-nique. The issue of detection limits becomes more importantif the objective of the study is to assess the quantity ofamorphous material present in what is a predominatelycrystalline sample because, as stated earlier, in a milledsample amorphous material may typically only be present upto 1% w/w. An assessment of the applicability of solutioncalorimetry to study small amorphous contents in solidpharmaceuticals was conducted by Hogan and Buckton [4],who prepared a calibration curve for lactose between 0 and10% w/w amorphous content in the same way as describedabove. They found that the technique could quantifyamorphous content to 0.5% w/w but noted that care neededto be taken when preparing the ampoules, because ingress ofeven small amounts of humidity caused partial recrystalli-sation of the sample before measurement.

Usually, in experiments designed to measure degrees ofcrystallinity of amorphous content a solvent is selected inwhich the solute is freely soluble. This ensures completedissolution of the sample within the time frame of theexperiment. Harjunen et al. [33] studied the dissolution of


lactose into saturated aqueous lactose solutions, a systemwhere clearly the solute would not completely dissolve.Interestingly, they observed a linear relationship between theamorphous content of the lactose solute and the measuredheat of solution in the saturated lactose solution (D satH).Similarly, a linear relationship was found between the amor-phous content of lactose and D satH in methanolic saturatedsolutions of lactose [34].

3.3. Characterisation of Interactions

As discussed above, application of Hesss law allows theconstruction of solution calorimetry experiments that allowvaluable information on the interaction of a solvent with amixed solvent system. An example would be to measure thedissolution of a solid in a buffer and a buffered solution ofmicelles. The difference in the D solH values recorded in thetwo systems would represent the enthalpy of transfer of thesolute from the buffer to the micelle. Similarly, if a complexsolvent system is employed, such as one that matched abiological medium for instance, it would be possible tomeasure the enthalpy of interaction of a solute with the entiremedium and then each of its constituents in turn.

This approach has been used to investigate the interactionof two solutes, propranolol HCl and mannitol, with twosimulated intestinal fluids (fasted-state and fed-state) andHanks balanced salt solution (HBSS) [35]. The simulatedintestinal fluids contained bile salts and lipids which formedmixed micelles while no micelles formed in HBSS. It wasfound that the two solutes exhibited endothermic heats ofsolution in all solvents; however, the values for propranololHCl were lower in the two simulated fluids than in HBSSwhile the values for mannitol were constant in all media.Calculation of the enthalpy of transfer (D transH) for pro-pranolol HCl into the micelles in the two simulated fluidsrevealed an exothermic (and hence favourable) interaction(-10.3 kJ mol-1 for the fed-state and -2.1 kJ mol-1 for thefasted-state).

In an earlier study, Beezer et al. [36] calculated values ofD transH for a series of alkoxyphenols to Escherichia coli cells,while the combination of Caco-2 cells and simulatedintestinal fluids has been suggested as a model for studyingdrug permeability through membranes [37]. It seems as ifsolution calorimetry may an ideal technique by which tomonitor these interactions, especially as it is unaffected bycloudy or turbid solutions or suspensions.

Chada et al. [38] used solution calorimetry to probe theinteractions of diclofenac sodium in cyclodextrin solutionsand water/ethanol mixtures. Tong et al. [39] used solutioncalorimetry to evaluate the stability constants and enthalpychanges associated with the formation of complexes between2-hydroxypropyl-b -cyclodextrin and a group of 12 aminecompounds which all had a diphenylmethyl functionalgroup. They found that only terfenadine HCl formed a 1:2complex with the b -cyclodextrin, the other 11 compounds allforming 1:1 complexes.

Solution calorimetry has also been used to measure theenthalpy of solution of diclofenac sodium, paracetamol andtheir binary mixtures [40] and to evaluate the in vitrocompatibility of amoxicillin/clavulanic acid and ampicillin/sulbactam with ciprofloxacin [41].

3.4. Other Applications

Perlovich and Bauer-Brandl [42] have discussed the useof heat of solution data to predict drug solubility using twomodel compounds, benzoic acid and aspirin. Their datasuggest it may be possible to predict the solubility orsolvation of a drug in different media. Similarly, Willson andSokoloski [43], as part of a study developing a method torank the stability of drug polymorphs, correlated solutioncalorimetry measurements to conventionally determinedsolubility data. Solution calorimetry has also been employedto study the dissolution and solvation of the model systems,flurbiprofen and diflunisal [44] and to compare the solvationof (+)-naproxen with three model NSAIDs (benzoic acid,diflunisal and flurbiprofen) [45].

Recent work has shown that solution calorimetry can beused to investigate the stability of supersaturated systems[46]. Supersaturated systems are particularly important fortopical and transdermal formulations where the API is mayformulated above its solubility in order to maximise thediffusional driving force for absorption [47, 48]. Clearly,supersaturated formulations are inherently thermodynami-cally unstable and it is hence likely that crystallisation of theactive will occur during storage with a resultant change indrug bioavailability. The ability to measure the time takenbefore recrystallisation occurs is therefore an essentialprerequisite to the development of such transdermal formula-tions and is difficult because the formulations are oftenopaque semi-solids.

Solution calorimetry allows the study of these formula-tions because it permits the in-situ formation of saturatedsolutions. Hadgraft et al. [46] have shown that by addingibuprofen in a co-solvent to a saturated solution of ibuprofenit is possible to form a supersaturated drug solution in thecalorimetric ampoule. In the absence of any stabilising com-pound the supersaturated system immediately precipitatesand a heat signal is observed. If a small concentration (0.1%w/v) of hydroxypropyl methyl cellulose (HPMC) is added tothe system the time to recrystallisation increased toapproximately 8 minutes. At a concentration of HPMC of0.5 % w/v the supersaturated system remained stable forapproximately 30 minutes, while at an HPMC concentrationof 1% w/v the system was stable for longer than 24 hours.

Barriocanal et al. [49] studied the formation of liposomeformation by coating the inside of a glass ampoule withphospholipids (deposited from a chloroform solution). Uponbreaking, the phospholipids film hydrated and liposomesformed; the solution calorimeter measured the changes inheat associated with the processes. They found that theformation of liposomes from egg phosphatidylcholine wasexothermic while the formation of liposomes from dimyris-toylphosphatidylcholine was endothermic and suggest thatthis difference arose from the influence of the hydrocarbonchains predominately on the hydration process. They alsonoted that the retention of small quantities of chloroform inthe phospholipids film significantly altered the enthalpychange of liposome formation, an effect that was ascribed tothe effect of chloroform on hydration.

Solution calorimetry can also be used to study the actionsof formulations directly. For instance, Gaisford et al. [50]


used solution calorimetry to follow the acid neutralisation ofmagnesium trisilicate mixture BP. They showed that of thethree active components in the mixture (magnesium carbo-nate, magnesium trisilicate and sodium bicarbonate), magne-sium carbonate contributed most to the action of the product,followed by sodium bicarbonate. The results suggested thatmagnesium trisilicate did not neutralise a significant quantityof acid by itself but that it extended the neutralising responseof the mixture, presumably by interacting with themagnesium carbonate. The results provided an insight intothe mechanism of action of the product and the data wouldbe useful in any reformulation exercise.

SUMMARY

Solution calorimetry offers the same virtues to thepharmaceutical scientist as other forms of calorimetry; it isnon-invasive, non-destructive and is not limited to the studyof homogeneous systems. It does not require large quantitiesof sample and the data are easy to interpret, so long as carehas been taken in experimental design and execution. Solu-tion calorimetry has classically been used during preformula-tion, where it allows the detection and identification ofpolymorphs and the quantification of crystalline or amor-phous contents. However, its ability to study cloudy or turbidsolutions or suspensions means it is finding increasingapplications in the direct study of the interactions of drugswith complex biological systems, where it is possible tomeasure the enthalpy of transfer of a drug between phases.Novel experimental strategies are also expanding its range ofapplications, and it has been used to investigate the stabilityof supersaturated systems, antacid formulations and theformation of liposomes. Its continued use and developmentwill ensure it retains its place as a vital tool in pharmaceut-ical formulation and preformulation.

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