Thermodynamics of aggregate formation between a non-ionic polymer and ionic surfactants: An isothermal titration calorimetric study

Accepted Manuscript Title: Thermodynamics of Aggregate Formation between a Non-Ionic Polymer and Ionic Surfactants: an Isothermal Titration Calorimetric Study Author: Salin Gupta Patel Paul M. Bummer PII: DOI: Reference:

S0378-5173(16)31020-1 http://dx.doi.org/doi:10.1016/j.ijpharm.2016.10.053 IJP 16184

To appear in:

International Journal of Pharmaceutics

Received date: Revised date: Accepted date:

20-6-2016 13-10-2016 23-10-2016

Please cite this article as: Patel, Salin Gupta, Bummer, Paul M., Thermodynamics of Aggregate Formation between a Non-Ionic Polymer and Ionic Surfactants: an Isothermal Titration Calorimetric Study.International Journal of Pharmaceutics http://dx.doi.org/10.1016/j.ijpharm.2016.10.053 This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

Thermodynamics of Aggregate Formation between a Non-Ionic Polymer and Ionic Surfactants: an Isothermal Titration Calorimetric Study

Salin Gupta Patel* and Paul M. Bummer

Department of Pharmaceutical Sciences College of Pharmacy University of Kentucky Lexington, KY 40536

*Corresponding Author Salin Gupta Patel University of Kentucky 789 South Limestone Street Lexington, KY 40536-0596 Email: [email protected]

Suggested running head: Aggregate Formation between HPMC and Ionic Surfactants: ITC Study

Page 1 of 52

ABSTRACT This report examines the energetics of aggregate formation between hydroxypropyl methylcellulose (HPMC) and model ionic surfactants including sodium dodecyl sulfate (SDS) at pharmaceutically relevant concentrations using the isothermal titration calorimetry (ITC) technique and a novel treatment of calorimetric data that accounts for the various species formed. The influence of molecular weight of HPMC, temperature and ionic strength of solution on the aggregate formation process was explored. The interaction between SDS and HPMC was determined to be an endothermic process and initiated at a critical aggregation concentration (CAC). The SDS-HPMC interactions were observed to be cooperative in nature and dependent on temperature and ionic strength of the solution. Molecular weight of HPMC significantly shifted the interaction parameters between HPMC and SDS such that at the highest molecular weight (HPMC K-100M; >240 kDa), although the general shape of the titration curve (enthalpogram) was observed to remain similar, the critical concentration parameters (CAC, polymer saturation concentration (Csat) and critical micelle concentration (CMC)) were significantly altered and shifted to lower concentrations of SDS. Ionic strength was also observed to influence the critical concentration parameters for the SDS-HPMC aggregation and decreased to lower SDS concentrations with increasing ionic strength for both anionic and cationic surfactant-HPMC ° ° systems. From these data, other thermodynamic parameters of aggregation such as 𝛥𝐻𝑎𝑔𝑔 , 𝛥𝐺𝑎𝑔𝑔 ° ° , 𝐻𝑎𝑔𝑔 , 𝛥𝑆𝑎𝑔𝑔 , and 𝛥𝐶𝑝 were calculated and utilized to postulate the hydrophobic nature of SDS-

HPMC aggregate formation. The type of ionic surfactant head group (anionic vs. cationic i.e., dodecyltrimethylammonium bromide (DTAB)) was found to influence the strength of HPMCPage 2 of 52

surfactant interactions wherein a distinct CAC signifying the strength of HPMC-DTAB interactions was not observed. The interpretation of the microcalorimetric data at different temperatures and ionic strengths while varying properties of polymer and surfactant was a very effective tool in investigating the nature and energetics of HPMC and ionic surfactant interactions.

Keywords: HPMC; ionic surfactants; isothermal titration calorimetry (ITC); polymersurfactant interactions, nanosuspensions.

INTRODUCTION Nanoparticle drug delivery systems (NDDS) are one of the several possible routes for bioavailability enhancement of poorly soluble drugs through enhanced dissolution rate, solubility, or both (1). However, NDDS often encounter varying degrees of thermodynamic instability leading to nanoparticle aggregation (1, 2). This instability is attributed to the extensive surface area that is generated by either of the two distinct approaches used to produce NDDS: (1) top down method where larger particles are broken down into nanoparticles through attrition, and (2) bottoms up method where new nanosized particles are created. The higher surface area is accompanied by a large positive free energy and without any effort to dampen the surface energy, the system prefers to move to an equilibrium state of the lowest free energy via aggregation of the smaller particles into larger particles (3-5).

Page 3 of 52

Nanosized particles generated by these approaches are typically physically stabilized by steric or ionic surface modification (6). Some recent reports have investigated the role of surface modifiers for providing physical stability to nanoparticles of various sizes and shapes(5). It has been observed that certain excipients such as surfactants and polymers are effective (i.e., steric or ionic surface modification) in producing physically stable NDDS by adsorption of free polymer, surfactant and polymer-surfactant aggregates on the surface of NDDS, potentially decreasing the surface energy of these particles (2, 7). Towards understanding the adsorption behavior of these polymer and surfactants, their concentration dependent interaction and speciation behavior needs to be explored. In general there are typically three critical concentration dependent behaviors reported for polymer-surfactant interactions (8).

The first critical concentration dependent

behavior is when the surfactant concentration exceeds a critical concentration where the surfactant begins to interact with the polymer, which is defined as a critical aggregation concentration (CAC) (8). The second is the polymer saturation concentration (Csat) when the surfactant molecules completely saturate the polymer chains and additional binding of the surfactant molecules cannot occur. This binding however, was observed by Persson et.al (9) to depend on the concentration of the polymer being studied and may be altered as a result of properties of the polymer such as concentration, molecular weight and polymer structure. The third concentration behavior is known as the critical micelle concentration (CMC) that results in the formation of surfactant micelles either after the saturation of the polymer in polymer-surfactant mixtures or pure surfactant micelles in the absence of a polymer (10). Goddard et. al(8) has provided a good review of these critical concentration parameters along with a surface tension/concentration plot of SDS in the presence

Page 4 of 52

of polyvinylpyrrolidone (PVP) at various concentrations; the various breakpoints (CAC, Csat and CMC) are shown with schematic assignments to direct the eye. Although there has been a good effort in the literature on understanding the interaction behavior of polymers and surfactant however, there is still a need to understand what parameters govern the energetics of the aggregation process of pharmaceutically relevant polymers such as PVP, HPMC and PEG as well as surfactants such as SDS and polysorbates at the concentrations typically used in formulation development. Moreover, some of the significant gaps identified from the review of these studies are: (1) the higher concentrations of HPMC typically utilized (3, 6, 11) in formulation of NDDS (0.25-1 %w/w) have not been evaluated, (2) the influence of molecular weight of polymer (HPMC) that could help provide insight into the number of binding site constraints for HPMC, (3) the effect of ionic strengths of solution and surfactant properties (i.e., head group, chain length) on the interaction between surfactants with HPMC has also not been investigated to the best of our knowledge, and (4)finally in our opinion the treatment of calorimetric data conducted while studying such polymer/surfactant systems (12) do not properly account for the enthalpies of dilution associated with the formation of various species (i.e., surfactant monomers, HPMC-surfactant aggregates, micelles) during the course of the titration experiment and needs to be treated more critically in order to allow for more accurate interpretation. This critical treatment of data will be discussed in detail in the methods section of this chapter. As discussed earlier since the effect of molecular weight, surfactant properties and ionic strength on the nature and strength of the HPMC-surfactant interactions remain unanswered, there is an additional need to understand which of these parameters govern the energetics of the aggregation process of pharmaceutically relevant polymers at the concentrations typically used in Page 5 of 52

drug product development and will be explored systematically in this work. Furthermore, at present there is very little molecular-level understanding that relates the interactions of surfactant and polymer in the bulk state to the adsorption of these excipients on solid surface that may ultimately lead to the final stability of NDDS. A lack of this mechanistic understanding and its relationship to the drug-excipient interactions in the bulk and the surface is a significant gap. This gap in the mechanistic understanding contributes to a complete lack in manipulation and control strategy for optimizing these excipients to maximize stability of NDDS. The development of this understanding can help in an a priori selection of type and level of pharmaceutical excipients to maximize the long term stability of NDDS with minimal experimentation. In order to develop this understanding, it would be critical to first develop an in-depth understanding of the concentration dependent speciation and energetics of the interactions between pharmaceutically relevant polymers and surfactants as well as how the excipient and solution properties influence the polymer-surfactant aggregate formation. The goal of this study was to determine the energetics of aggregate formation between a model non-ionic polymer, hydroxypropyl methylcellulose (HPMC),

and

model

ionic

surfactants,

sodium

dodecyl

sulfate

(SDS),

hexadecyltrimethylammonium bromide (CTAB), and dodecyltrimethylammonium bromide DTAB. The influence of ionic strength, molecular weight of HPMC, and type of ionic surfactant head group on the aggregate formation process was explored using an isothermal titration calorimetry (ITC) method. ITC technique has gained momentum over the last couple of decades to investigate the thermodynamics of micellization where researchers have determined critical micelle concentration (CMC) parameters, enthalpy, Gibbs free energy and entropy of micelle formation. In recent years, Page 6 of 52

the use of ITC for understanding polymer-surfactant interactions has grown particularly due to the increased sensitivity of the technique which has been utilized to quantitatively study the aggregation process between polymers and surfactants including polyethylene glycol (PEG)/SDS (13), ethyl hydroxyethyl cellulose (EHEC)/HPMC/SDS (19), polyethylene oxide (PEO)/SDS (1418), and polyvinylpyrrolidone (PVP)/sodium dodecylbenzenesulphonate (SDBS) (12, 13, 16, 17, 19, 20) systems. A number of good reviews are available for more detailed description and application of ITC (13, 21, 22). The ITC method was utilized to conduct a full thermodynamic characterization of the aggregation parameters between HPMC and the model ionic surfactants since it provides quantitative information on the critical concentration parameters (i.e., CAC, Csat and CMC) and energetics of the aggregation process (16). The critical thermodynamic parameters derived from ITC enthalpograms are known to contain information about rearrangement of polymers, micelle/aggregate formation and their disassociation, however, the interpretation of enthalpograms is still not very well developed area and hence there is significant scope for further understanding (15). The fundamental information can be obtained on how molecular properties of HPMC and ionic surfactants influence these parameters that may provide further insights into the molecular interactions driving the process. The learning from this study would be applied to future studies in characterizing the nature of HPMC-ionic surfactant aggregates and their adsorption onto the surface of nanoparticles.

MATERIALS AND METHODS Materials

Page 7 of 52

Sodium dodecyl sulfate (>98%) was obtained from Sigma-Aldrich (St. Louis, MO). The SDS thus obtained was further purified by the solid phase extraction process by passing a 1%w/w aqueous solution of SDS through a Waters SEP-PAK® plus C18 environmental cartridge, the extracted solution was then lyophilized. Surface tension measurements were conducted for lyophilized SDS to assess the presence of any local minimum near the critical micelle concentration (CMC) of the surfactant by the du Noüy ring tensiometer. An absence of such a local minimum allowed the usage of the purified surfactant. Purified and unpurified SDS solutions were however not found to be affected by the presence of trace quantities of dodecanol that are commonly found in the commercially available SDS. Hence the SDS purchased from SigmaAldrich (St. Louis, MO) was used as received. The cationic surfactants including DTAB, and, hexadecyltrimethylammonium bromide (CTAB) were purified in the same way as SDS with the Waters SEP-PAK® plus C18 environmental cartridges and the extracted solution was then lyophilized and examined for the presence of any local minimum near the CMC. The absence of a local minimum resulted in using these chemicals as received from Sigma-Aldrich (St. Louis, MO). Hydroxypropyl methylcellulose (Benecel® K-4M, K-15M and K-100M) were obtained from Ashland Aqualon Functional Ingredients, Ashland Inc (Wilmington, DE) and used as received. Sodium Chloride (NaCl) was purchased from Sigma-Aldrich (St. Louis, MO) and used as received. All solutions were prepared using purified water (18.2 MΩ cm) obtained from a MilliQ water purification system (Millipore, Billerica, MA). Methods The calorimetry experiments were performed using a TAM III isothermal titration calorimeter (ITC) manufactured by TA Instruments (New Castle, DE). TAM III calorimeter works Page 8 of 52

in power compensation mode (Fig.1) where the temperature of the sample cell is maintained constant using a temperature sensor with a feedback system utilizing a reference cell. When an enthalpic event (i.e., chemical reaction, molecular reorganization, binding, solubilization) occurs during titration which is either endothermic or exothermic, power is supplied to a the heater or cooler to bring the system back to isothermal state and was directly measured (23). For the titration, as per the experimental conditions, surfactant solutions of known concentrations above the CMC

(~10x CMC; 87.3mM for SDS, 129.7mM for DTAB and 9.1mM for CTAB,

respectively) were prepared and loaded into either a 1 ml or 5ml syringe mounted on a precision pump. In all experiments, a micellar surfactant solution of a model surfactant (~10x CMC) is added from a syringe in a step-wise manner into a thermodynamically isolated containing a known quantity of either deionized purified water (2.1 ml or 2.4 ml) or an HPMC solution (0.25 and 0.5% w/w basis) at various ionic strengths as per the experimental requirements at a selected temperature. At the start of each experiment the reference cell contained the same solution (i.e., similar physical and thermal properties) as the sample cell. Titrations at each predetermined time interval were performed by the syringe under computer control that injected either 20 or 25 µl of the surfactant solution into the sample cell. Usually 5 to 7 minutes were provided between injections to allow time for thermal equilibration while the sample cell was continuously stirred (60 rpm) with a turbine stirrer. Thus in a given ITC measurement, energetics associated with a process was directly measured at a constant temperature; for example if an enthalpic event that was exothermic occurred during titration, the compensation mode (feedback loop) would stop heating the sample cell and then cooled until the temperature of the sample cell was brought to the temperature of the reference cell. The signal (peak) thus obtained from the feedback loop was

Page 9 of 52

integrated directly to yield the heat associated (Q) with the event that occurred as a result of the given titration. As shown in Fig. 2a, the heat signal (Q) was directly related to the concentration and volume of the titrant in each injection which was normalized with respect to moles (δn) of surfactant added to the sample cell. It was further analyzed to obtain the apparent enthalpy (Q/δn) change using TAM III Lab Assistant Software provided by TA Instruments. The apparent enthalpy(ΔHapp ) was plotted against surfactant concentration (i.e. enthalpograms) (Fig. 2b). Enthalpogram analysis was based on the pattern formed by the plots as a result of the injections and the total composition in the sample cell. The number of titrations and experimental conditions were varied depending on the factors being studied in the experiment. All experiments were repeated at least twice to check for reproducibility of the measurements and for reference the representative SDS-HPMC interaction data are shown in Appendix.1. RESULTS AND DISCUSSION The goal of this study was to determine the effect of various parameters including the presence of a model non-ionic polymer (HPMC), temperature, ionic strength, and excipient properties (i.e., molecular weight of polymer, head-group and chain length of surfactant) on the energetics of the HPMC-surfactant aggregation process using ITC. ITC Method Validation using Ionic Surfactant Micellization Parameters The properties of ionic surfactant micelles (without HPMC) have been studied previously utilizing techniques such as ITC, NMR and surface tensiometry (24-26). Hence we first validated our ITC methodology by comparing the values of micellization parameters (i.e., CMC, enthalpy,

Page 10 of 52

entropy, and free energy of micellization) determined in the present study with those reported in the literature for three model ionic surfactants, SDS, CTAB, and DTAB (27-29). In all experiments, the apparent enthalpy 𝛥𝐻𝑎𝑝𝑝 determined from the heat signal measured by the ITC instrument (see “Methods” for more details) was plotted against surfactant concentration (i.e. enthalpograms). The enthalpograms (𝛥𝐻𝑎𝑝𝑝 vs. surfactant concentration) for CTAB, SDS, and DTAB are presented in Fig. 2b, 3a, and 3b. As the micellar surfactant solution is mixed with deionized water in the sample cell, an extremely dilute solution of surfactant monomers (demicellization) is formed. In Figs. 2b and 3, the enthalpograms can be subdivided within two concentration ranges where the enthalpies of reaction are almost constant. The first plateau seen in the enthalpogram is denoted as the demicellization plateau, which could be attributed to the cumulative heat changes due to the demicellization/breakdown of the surfactant micelles, as the micellar surfactant solution present in the syringe is being diluted in the sample cell in the initial titrations. These reaction enthalpies may also be accompanied along with the enthalpies associated with the dilution of surfactant monomers, and interactions with their counterions (30, 31). With subsequent titrations, as the surfactant concentration in the sample cell reaches the critical micelle concentration (CMC) of the surfactant, a sharp decrease in the apparent enthalpy is observed at a concentration of ~8mM for SDS, ~14mM for DTAB, and ~1mM for CTAB, respectively. After this point the titration process mainly involves titrating the micellar surfactant solution in the syringe into the micellar solution already present in the sample cell (second plateau in the enthalpogram; Fig. 2b and 3). This range suggests that the enthalpies associated with the dilution of micelles is being measured. Similar ITC profiles have been reported for ionic surfactants anionic surfactants such as SDS, SDES (sodium decylsulfate), SDBS and Page 11 of 52

cationic surfactants such as CTAB, DTAB and TTAB. Blandamer et al.(22) described a similar shape of the ITC plot for SDS and SDES micelle deaggregation. The consistent view in the literature describes that the change in apparent enthalpy for surfactant micellization may include various components such as shown below in equation (1). ° The enthalpy of demicellization (𝛥𝐻𝑑𝑒𝑚𝑖𝑐 ) is equal in magnitude to the enthalpy of micellization

although it bears an opposite sign (23, 29, 32). ° 𝛥𝐻𝑎𝑝𝑝 = 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) + 𝛥𝐻𝑚𝑖𝑐 + 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠)

(1)

where 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) and 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) are the enthalpies of dilution of the surfactant ° monomers and surfactant micelles, respectively. 𝛥𝐻𝑚𝑖𝑐 represents the standard enthalpy of

micellization and expressed per mole of surfactant monomer unit.

𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) and

𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) were measured independently by direct experimentation utilizing ITC (i.e., by titrating a dilute surfactant solution and a micellar surfactant solution into the sample cell containing the same dilute and micellar surfactant solution present in the syringe, respectively). Although the independent titrations to measure 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) and 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) were conducted, they were ° determined to be negligible for all the three model systems and hence 𝛥𝐻𝑚𝑖𝑐 was directly obtained

from enthalpograms (𝛥𝐻𝑎𝑝𝑝 vs. surfactant concentration). To determine the standard enthalpy of ° micellization (𝛥𝐻𝑚𝑖𝑐 ) directly from the enthalpogram, lines are drawn to fit the two plateaus above

and below the observed inflection point i.e., CMC. A line is drawn perpendicular to the x axis, ° and 𝛥𝐻𝑚𝑖𝑐 is measured from the length of the segment connecting the two extrapolated lines (Fig. ° 2b). The 𝛥𝐻𝑚𝑖𝑐 for SDS at 25 ˚C was determined to be -0.7 kJ/mol (Table 1), which is close to

the value of -0.75 kJ/mol reported by Woolley et al.(27) and -0.5 kJ/mol by Singh et al.(12). In Page 12 of 52

° case of DTAB, the 𝛥𝐻𝑚𝑖𝑐 at 25 ˚C was determined to 1.9 kJ/mol (Table 2), which is similar to the

value of 1.87 kJ/mol reported by Beyer et al (29). The CMC value is determined from the extremum (highest peak) of the first derivative of the 𝛥𝐻𝑎𝑝𝑝 vs. surfactant concentration plot (Fig. 2c). For SDS, the CMC was determined to be 8.3 mM (Table 1) and is in agreement with the values reported by Philips et al. (8.1 mM) and Horin et al. (8.2 mM) (33-35). The CMC for CTAB and DTAB at 25°C was determined to be 1.25 mM and 14.49 mM, respectively (Fig. 2b and Table 2). Blandamer et al.(36) reported the CMC for CTAB at 25°C to be 0.97 mM utilizing ITC while Beyer et al.(29) utilizing the same technique reported the CMC for DTAB at 25°C as 13.5 mM, both of which are similar to the determined ° values here. Overall, the micellization parameters such as CMC and 𝐻𝑚𝑖𝑐 for three model ionic

surfactants determined in this study were in good agreement with the literature values, which validated our ITC methodology. Thermodynamics of Model Ionic Surfactant Micellization: Phase Separation Model ° Additional thermodynamic parameters such as standard free energy (𝛥𝐺𝑚𝑖𝑐 ), standard ° ° enthalpy (𝐻𝑚𝑖𝑐 ), standard entropy (𝛥𝑆𝑚𝑖𝑐 ), and heat capacity (𝛥𝐶𝑝 ), respectively of micellization

of model ionic surfactants SDS and DTAB were determined at different temperatures using the phase separation model (37). According to the phase separation model, the micellization behavior is not assumed to be a progressive stepwise association of surfactant monomers but rather the result of a one-step process leading to the formation of a separate phase in the bulk aqueous phase. In other words, at or above CMC the system contains two separate phases (i.e., surfactant monomers and surfactant micelles) that are in equilibrium. When the micelle aggregation number is large,

Page 13 of 52

the phase separation model has been shown to be special case of the mass action model in which micellization is considered a stepwise association of monomers(38). The micellization process can be represented by equation (2). 𝑛𝑆 − + 𝑚𝑋 + ⇔ [𝑆𝑛 𝑋𝑚 ](𝑛−𝑚)−

(2)

𝐾

where each micelle ([𝑆𝑛 𝑋𝑚 ](𝑛−𝑚)− ) formed is assumed to contain 𝑛 surfactant ions (𝑆 − ) and 𝑚 dissociated counterions (𝑋 + ) with the fraction of charge for surfactant ions of 𝑚/𝑛 (degree of ionization; 𝛼) in each micelle to give electroneutrality. The equilibrium constant (𝐾) can be expressed as equation (3) [𝑆𝑛 𝑋𝑚 ](𝑛−𝑚)− 𝐾= [𝑆 − ]𝑛 [𝑋 + ]𝑚

(3)

° The standard free energy of micellization (𝛥𝐺𝑚𝑖𝑐 ), which is the standard free energy per

mole of surfactant monomer unit, 𝛥𝐺 ° ⁄𝑛 , is ° 𝛥𝐺𝑚𝑖𝑐

𝛥𝐺 ° 𝑅𝑇 = = 𝑛 𝑛 ln 𝐾

(4)

where 𝑛 is the micelle aggregation number. As per the phase separation model, at CMC, [𝑆 − ]= [𝑋 + ]=CMC (27, 39). Thus, at CMC equation (4) can be computed as 1

° 𝛥𝐺𝑚𝑖𝑐 = −𝑅𝑇 [𝑛 ln[𝑆𝑛 𝑋𝑚 ](𝑛−𝑚)− − ln(𝐶𝑀𝐶) −

𝑚 𝑛

ln(𝐶𝑀𝐶)]

(5)

In the evaluation of the thermodynamic parameters, the first term in equation 5 can be neglected due to their low weightage on the parameters and equation (5) is computed as (30) ° 𝛥𝐺𝑚𝑖𝑐 = 𝑅𝑇 (1 + 𝛼) ln(𝐶𝑀𝐶)

Page 14 of 52

(6)

where

𝑚 𝑛

used as α is the counterion association for the micelle and is assumed to be one

by accounting for 100% counterion binding on the micelle for ionic surfactants (8) and thus equation 6 becomes ° 𝛥𝐺𝑚𝑖𝑐 = 2 𝑅𝑇 ln(𝐶𝑀𝐶)

(7)

In equation (7) it is assumed that the ionic micelle has net electroneutrality. This is attributed to the self-aggregation of the monomers of the ionic surfactant and their binding with an equal number of counterions to form neutral micelles (30). The effect of temperature on the CMC ° and 𝛥𝐻𝑚𝑖𝑐 for SDS and DTAB was determined at 25°C, 32°C, 40°C and 50°C (Figs. 3 and Table ° ° 1-2). The thermodynamic parameters such as 𝛥𝐺𝑚𝑖𝑐 and 𝛥𝑆𝑚𝑖𝑐 for SDS and DTAB were obtained ° from equations (7) and (8) (Tables 1 and 2). The 𝛥𝐺𝑚𝑖𝑐 is twice that of a nonionic micelle since

two species bind together to form a micelle for ionic surfactants. ° ° ° 𝛥𝐺𝑚𝑖𝑐 = 𝛥𝐻𝑚𝑖𝑐 − T𝛥𝑆𝑚𝑖𝑐

(8)

° The 𝛥𝐺𝑚𝑖𝑐 of SDS and DTAB micellization at 25°C was also determined to be strongly ° negative (spontaneous process) and remained favorable at higher temperatures. The 𝛥𝑆𝑚𝑖𝑐 of SDS ° and DTAB micellization was observed to decrease and while 𝛥𝐻𝑚𝑖𝑐 increased at higher

temperatures and due to an enthalpic-entropic compensation, the overall free energy of the micellization process remained fairly constant (29, 32). Volpe and Wang et al. have also determined the SDS micellization process to be enthalpically unfavorable and hence have suggested it to be entropically driven at 25°C. Similar to these observations, in this work the SDS micellization process is shown to be more enthalpically favored with increasing temperatures (14, 40). The CMC of both SDS and DTAB shifted slightly with increasing temperature indicating that Page 15 of 52

the micellization of SDS and DTAB is mainly driven by hydrophobic interactions (41-43). Another thermodynamic parameter is the heat capacity of micellization (𝛥𝐶𝑝 ) that is obtained from ° the slope of 𝛥𝐻𝑚𝑖𝑐 vs 𝑇. 𝛥𝐶𝑝 is -0.5 kJ/mol and -0.4 kJ/mol for SDS and DTAB micellization,

respectively (Tables 1 and 2). The strongly negative value of 𝛥𝐶𝑝 for SDS and DTAB micelles is indicative of the transfer of SDS/DTAB monomers from the aqueous phase into the more hydrophobic micellar phase (36, 44-46). These results are in agreement with the consistent view in the literature where hydrophobic effect is considered the predominant driving force for micellization of ionic surfactants (27, 46). Influence of HPMC on the Energetics of HPMC-Ionic Surfactant Aggregation Process To investigate the energetics of interactions between HPMC and model ionic surfactants (i.e., SDS and DTAB) and to interpret plausible mechanisms, ITC and the phase separation model (see previous section for more detail) were utilized. Fig. 4a shows the enthalpograms for titration of an SDS micellar solution (10x CMC; ~87.3mM) into the sample cell containing 0.25%w/w or 0.5% w/w HPMC K-4M solution at 25°C. A distinct endothermic peak is observed upon titration of SDS into HPMC K-4M where it is evident that the presence of HPMC changes the shape of the enthalpograms, deviating from the SDS/water enthalpograms (Figs. 3 & 4a). This difference between the enthalpograms for the SDS/HPMC/water and SDS/water systems could be attributed to the interactions between SDS and HPMC (15). In Fig 4, similar to micellization of pure ionic surfactants (previous section), the change in apparent enthalpy for polymer-surfactant interactions (𝛥𝐻𝑎𝑝𝑝 ) may consist of various components as shown below in equations (9) and (10) (23, 32, 44).

Page 16 of 52

° 𝛥𝐻𝑎𝑝𝑝 = 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) + 𝛥𝐻(𝑝−𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝−𝑠) + 𝛥𝐻𝑚𝑖𝑐 + 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝)

(9)

where 𝛥𝐻𝑐𝑜𝑟𝑟 is the corrected enthalpy of interaction for SDS-HPMC systems, 𝛥𝐻𝑑𝑖𝑙(𝑝) is the enthalpy of dilution of the polymer, 𝛥𝐻(𝑝−𝑠) is the enthalpy of polymer-surfactant interaction and 𝛥𝐻𝑑𝑖𝑙(𝑝−𝑠) is the heat of dilution of the HPMC-surfactant aggregates. The treatment of the raw data obtained from the microcalorimeter in the literature follows ° either the direct interpretation of the 𝛥𝐻𝑎𝑝𝑝 (16, 17) or subtraction of the entire surfactant

demicellization titration curves in the absence of polymer (12). Although a correction of the raw microcalorimetric data is required for the accurate interpretation of data, it is our opinion that subtraction of the entire surfactant demicellization (also known as dilution) curve ° (i.e., 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) + 𝛥𝐻𝑚𝑖𝑐 + 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) ) from the SDS/HPMC enthalpogram may not be

suitable in all the regions of the enthalpogram. This may be unsuitable since it is known that the presence of HPMC significantly alters the enthalpy of dilution dependence on SDS concentration associated with SDS monomeric fraction. Thus, a direct subtraction of the enthalpy of dilution (associated with titration of SDS into water) is expected to overcorrect especially in the region associated with the micellization of SDS which occurs at lower SDS concentrations in the SDS/water system as compared to significantly higher SDS concentrations for the SDS/HPMC/water systems. Similarly, an under correction could apply to the region of the enthalpogram where the free micelles begin to occur in the SDS/HPMC/water system. A simple subtraction of the SDS dilution enthalpogram without consideration of the species present in the experimental conditions could potentially result in inaccurate values of the thermodynamic parameters associated with the aggregation and micellization of the SDS-HPMC and free SDS

Page 17 of 52

micelles, respectively. Thus, for the purposes of this work and a more accurate interpretation of the calorimetric data, the novel approach to the correction of 𝛥𝐻𝑎𝑝𝑝 will be conducted as per the species expected at equilibrium in the sample cell at each given titration. In other words for the titrations where SDS monomers are expected to predominantly occur in the sample cell (i.e., <17 mM SDS concentration at 0.25% w/w HPMC and 25°C or as described later region A→C) the 𝛥𝐻𝑐𝑜𝑟𝑟 value obtained is processed as per equation (10). Similarly, at the SDS concentrations where the formation of free SDS micelles are expected during the titration experiment (i.e., 21-24 mM SDS concentration at 0.25% w/w HPMC and 25°C or as described later region C→D), the correction (𝛥𝐻𝑐𝑜𝑟𝑟 ) applied is shown in equation (11) and finally at SDS concentrations > 24mM or region D where it is predominantly all of the SDS being titrated are expected to form SDS free micelles the correction (𝛥𝐻𝑐𝑜𝑟𝑟 ) applied is shown in equation (12). This approach should address the over and under correction of the raw data mentioned above. 𝛥𝐻𝑐𝑜𝑟𝑟 = 𝛥𝐻(𝑝−𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝−𝑠) = 𝛥𝐻𝑎𝑝𝑝 − ( 𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝) )

(10)

° 𝛥𝐻𝑐𝑜𝑟𝑟 = 𝛥𝐻(𝑝−𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝−𝑠) = 𝛥𝐻𝑎𝑝𝑝 − ( 𝛥𝐻𝑚𝑖𝑐 + 𝛥𝐻𝑑𝑖𝑙(𝑝) )

(11)

𝛥𝐻𝑐𝑜𝑟𝑟 = 𝛥𝐻(𝑝−𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝−𝑠) = 𝛥𝐻𝑎𝑝𝑝 − ( 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) + 𝛥𝐻𝑑𝑖𝑙(𝑝) )

(12)

For the 𝛥𝐻𝑐𝑜𝑟𝑟 determination, the various heats of dilution were measured independently by direct experimentation utilizing the ITC and were subtracted as per equations (10, 11 and 12). Additionally, to minimize the influence of enthalpies of dilution shown above, the temperature of 25°C was selected to study SDS/HPMC interactions because at this temperature the total ° contributions from the dilution of SDS (𝛥𝐻𝑑𝑖𝑙(𝑚𝑜𝑛,𝑠) − 𝛥𝐻𝑚𝑖𝑐 + 𝛥𝐻𝑑𝑖𝑙(𝑚𝑖𝑐,𝑠) ) were determined to

be close to zero. Fig. 4a shows the enthalpograms at 25°C for pure SDS in water (0% HPMC) and Page 18 of 52

SDS/HPMC system (0.25% and 0.5% HPMC), whereas Fig. 4b shows the corrected enthalpy (𝛥𝐻𝑐𝑜𝑟𝑟 ) plotted as function of SDS concentration at 25°C where the corrected enthalpograms preserve the shape of the uncorrected enthalpograms lending further to the selection of 25°C as the ideal temperature for investigating the SDS-HPMC interactions. For simplification after the initial few titrations at low SDS concentrations where little to no interactions are seen, the SDS/HPMC/water enthalpogram (Fig. 4b) can be divided into three main regions: (1) A→B denotes the sharp increase in the enthalpy (endothermic region) reaching a maximum, (2) B→C depicts decrease in enthalpy to a minimum (exothermic region), and finally (3) C→D represents an increase in enthalpy reaching a plateau. 1. SDS-HPMC Enthalpogram: A→B region In the A→B region (Fig. 4b), a sharp increase in the enthalpy at ~4 mM SDS concentrations shows the first critical concentration denoted as the critical aggregation concentration (CAC). As shown in Fig. 4b, when the SDS concentration increases above the CAC, an endothermic maximum is observed indicating increased interactions between HPMC and SDS. The sudden increase in enthalpy in this region could signal cooperative interactions between SDS and HPMC at 25°C. Goddard et al. (8) provided a review of early studies describing the presence of a break point (CAC) and suggested that it represented the start of interactions for polymer and surfactant interactions. In other words, greater SDS-HPMC interactions further facilitates the interaction between SDS and HPMC and the sharp and sudden increase in SDS-HPMC interactions suggests that they are cooperative in nature (47). The cooperative interactions between SDS and HPMC could be expected due to the availability of multiple sites for interactions between SDS and HPMC chains. Generally, interactions are considered as cooperative when the binding of one molecule Page 19 of 52

(or ligand) at one site affects the binding of the other molecules at other binding sites of the macromolecule (44). The affinity of ligand for the binding sites in such cases has been shown to be influenced by the amount of ligand that is already bound (48). Moreover, the CAC was also observed to be influenced by polymer concentration with a slightly lower value of CAC and sharper slope of the endothermic curve at the higher concentration of HPMC (0.5% w/w) reflecting an increase in SDS/HPMC cooperative interactions (Fig. 4b & Table 3). While Dai et al. confirmed interactions between PEG & SDS through a distinct CAC measured by a conductometry technique, the effect of PEG concentration on the CAC was not observed(16). Olofsson et al. (8, 13) using microcalorimetric studies reported that the CAC for surfactants and uncharged polymers in aqueous solutions was weakly dependent on the concentration of the polymer. The lack of the polymer concentration effect on CAC in some abovementioned studies may be related to the sensitivity of the method used to determine CAC. For example, the microcalorimetric method may show the polymer concentration effect as compared to the conductometry method due to higher sensitivity of the former method. Nilsson (49) used an equilibrium dialysis method and showed that the amount of SDS adsorbed onto HPMC increased linearly to a maximum beyond CAC of ~4.4 mM SDS. The author also measured the aggregation number of SDS aggregates adsorbed onto HPMC using a fluorescence quenching technique, which showed that the aggregation number increased from ~10 at CAC to a maximum of ~50 as the SDS concentration increased beyond CAC. The degree of cooperativity was higher at 0.2% HPMC as compared to 0.05% HPMC. Based on the intrinsic viscosity and Huggins constant (a measure of polymer-polymer interaction), a critical polymer concentration (~0.15%) was identified above which the intrinsic viscosity values increased linearly as the SDS concentration increased beyond Page 20 of 52

CAC. This was attributed to an intermolecular aggregation/clustering process creating a more extended three-dimensional network of HPMC-SDS aggregates.

In our experiments, the

concentration of HPMC is selected above this critical concentration (i.e. 0.25-0.5% w/w) and hence intermolecular aggregates are expected(49). The SDS and HPMC could exhibit hydrophobic and hydrogen bonding interactions since it is known that along with being moderately hydrophobic in nature, HPMC also has both hydrogen bond acceptor and donor groups and SDS head groups has hydrogen bond acceptor groups. Therefore, the question in this case would be which of the two the driving forces is dominant in our studies. Considering the moderately hydrophobic nature of HPMC, it was postulated that these interactions might be driven by the hydrophobic effect. In order to support this hypothesis, the influence of temperature on the CAC and the endothermic peak of the SDS-HPMC interactions was investigated. It has been shown in previous studies that if the polymer-surfactant interactions are driven by the hydrophobic effect, with increasing temperature CAC and the endothermic peak diminishes, which was mainly attributed to easier and more continuous type of interactions due to the breakdown of the hydration shell (water structure) surrounding the hydrophobic regions of polymers and surfactants at higher temperatures (28, 45) (50) (31). Dill et al.(50) described the hydrophobic effect as having a characteristic temperature dependence in the transfer of a nonpolar solute into an aqueous phase. In Fig. 5, with increasing temperatures while the general shape of the titration curves remain similar, the CAC is significantly altered such that as the temperature increases from 25°C to 40°C, the CAC is no longer as pronounced along with a prominent decrease in the endothermic peak with the CAC being almost undistinguishable at 40°C. The temperature dependence of CAC and the endothermic peak for HPMC-SDS system supports the hypothesis Page 21 of 52

that the driving force for the HPMC and SDS interactions is characteristic of the hydrophobic effect. Furthermore, it is known that the hydrophobic regions of HPMC consist of the methyl and hydroxypropyl groups (51). While the enthalpy data for the transfer of a hydroxypropyl group from aqueous to more hydrophobic phases are not readily available in the literature, the transfer enthalpy of a similar functional group, propan-1-ol, has been reported to be approximately +10 kJ/mol at 25°C and the transfer enthalpy for a methyl group is zero (52). This further suggests that the interactions would thus be endothermic and enthalpically unfavorable. 2. SDS-HPMC Enthalpogram: B→C region In the next B→C region (Fig. 4b) of the titration curve as the SDS concentration increases further, a sharp almost linear decrease in the endothermic peak is seen leading to an exothermic minimum at ~17 mM SDS concentration (0.25% HPMC) at 25°C. This decrease in the apparent enthalpy could be indicative of decreased hydrophobic interactions. To test this hypothesis the effect of temperature in this region was determined. As shown in Fig.5, the enthalpograms vary with temperature suggesting that hydrophobic interactions could be involved in the B→C region. However, the enthalpy change in this region becomes more endothermic with an increase in temperature, which is opposite of what is generally expected with hydrophobic interactions. This inverse temperature dependence may suggest a strong decrease in the hydrophobic interactions. Nilsson (49) showed that the SDS monomer concentration remained constant at ~4 mM when the total SDS concentration increased from ~5 mM to ~16 mM (0.2% w/w HPMC). As the SDS concentration increased, the SDS-HPMC aggregate grew as reflected by linear increase in the aggregation number from <10 to 50. The intrinsic viscosity and the cloud point of the SDS-HPMC system steadily decreased and increased, respectively, when SDS concentration increased from ~7 Page 22 of 52

mM to ~12 mM and reached a plateau beyond 12 mM SDS concentration. This was attributed to the restructuring of HPMC-SDS aggregate network, where the intermolecular networking capability of HPMC is reduced resulting in lower viscosity(49).

The restructuring HPMC-SDS

aggregate network at higher SDS concentrations could be attributed to the expansion and rehydration of the HPMC chains caused by the electrostatic repulsion exerted by the anionic head groups of the densely adsorbed SDS (43, 44). Moreover, as discussed earlier, the rehydration of hydrophobic groups of HPMC such as hydroxypropyl and methyl groups is expected to be exothermic based on the reported enthalpy of hydration of -10 kJ/mol for propan-1-ol (52) and 0 kJ/mol for methyl (53) groups at 25°C. Thus, as shown in Fig.4, the decrease in endothermic peak and the exothermic nature of the curve may be attributed to restructuring of HPMC-SDS aggregates upon higher adsorption of SDS on the HPMC chains (19, 41, 54). Overall, the temperature effect and the enthalpy of hydration support the hypothesis of decreased hydrophobic interactions between SDS and HPMC in the B→C region. Torn et al. (19) reported an inverse temperature dependence of the PVP-sodium dodecylbenzenesulphonate (SDBS) interactions in the exothermic region. The authors proposed a pearl-necklace model for the formation of PVP- SDBS aggregates. It was suggested that an expansion of the polymer chains due to the electrostatic repulsion of the adsorbed surfactant molecules on the polymer surface could allow the polymer chain to wrap around the surfactant aggregates. As shown in Fig. 4b, the exothermic minima (i.e., 17 mM for 0.25% HPMC at 25°C) may indicate saturation of HPMC chains with the adsorbed SDS, which is also known as the polymer saturation concentration (Csat)(8, 49). The exothermic minima shifts from 17 mM at 0.25% HPMC to 20 mM at 0.5% HPMC concentration, which supports the fact that more SDS is needed to saturate a greater number of HPMC binding sites.

Page 23 of 52

Nilsson(49) showed with equilibrium dialysis for 0.2% w/w HPMC at 25°C, all excess surfactant above the CAC is adsorbed onto the HPMC until a polymer saturation concentration (Csat~17 mM), followed by the an increase in free SDS monomer concentration until the formation of pure SDS micelles at CMC(~20 mM).

Moreover, with the aggregation numbers of the SDS-HPMC

aggregates the author showed that the overall number of aggregates increased along with an increase in the number of binding sites on HPMC with an increases in SDS concentration at this HPMC concentration. A model was proposed by Dai et al. (55) for PEG-SDS aggregate formation where the enthalpies of interaction between SDS and PEG were interpreted to be the result of polymer induced aggregation at the lower SDS concentration and rehydration of PEG chains at the higher surfactant concentrations. 3. SDS-HPMC Enthalpogram: C→D region Finally in the last C→D region of the titration curve, the enthalpy increases somewhat with increasing SDS concentrations and reaches a plateau at approximately zero enthalpy (Fig. 4b). This may be signaling that after the available binding sites on HPMC are saturated with SDS, the newly added SDS monomers mutually begin to interact until a critical micelle concentration (Cm) for SDS is achieved and pure SDS micelles begin to form. Nilsson (49), at 0.2% HPMC and 20°C, showed that at a total SDS concentration of ~17 mM, the free SDS monomer concentration start increasing and reach the Cm at ~20 mM. It was suggested that the HPMC binding sites were saturated with SDS at Csat of ~17 mM and hence any newly added SDS remained in the monomeric state until Cm was achieved. Above Cm, it is assumed that SDS-HPMC aggregates and pure SDS micelles coexist. As shown in Fig. 4b, the Cm for SDS were 21mM and 25 mM for 0.25% HPMC and 0.5% HPMC, respectively. The enthalpy change in this region for SDS-HPMC system is Page 24 of 52

approximately zero and superimposes with the enthalpy of pure SDS micelles dilution region in the enthalpograms and was least influenced by temperature, which supports the formation of pure SDS micelles above Cm. Thermodynamic Parameters of HPMC-SDS Aggregation For a more in depth understanding of the driving forces for the SDS-HPMC interactions, ° ° the standard free energy of aggregation (𝛥𝐺𝑎𝑔𝑔 ), standard enthalpy of aggregation(𝐻𝑎𝑔𝑔 ), standard ° entropy of aggregation(𝛥𝑆𝑎𝑔𝑔 ), and heat capacity (𝛥𝐶𝑝 ) of aggregation are determined (Table 3). ° The 𝛥𝐺𝑎𝑔𝑔 was computed using the phase separation (37) model described in the previous section

(equation (7)). The standard free energy of aggregation per mole of surfactant monomer unit, 𝛥𝐺 ° ⁄𝑛 , is ° ⁄ 𝛥𝐺𝑎𝑔𝑔 𝑛 = 𝑅𝑇⁄𝑛 ln 𝐾

(13)

where 𝑛 is the number of moles of surfactant and [𝑆 − ]= [𝑋 + ]=CAC; thus at CAC it can be computed as equation (14) that can be further rearranged to equation (15). ° 𝛥𝐺𝑎𝑔𝑔 = 2 𝑅𝑇 ln(𝐶𝐴𝐶)

(14)

° ° ° 𝛥𝐺𝑎𝑔𝑔 = 𝛥𝐻𝑎𝑔𝑔 − T𝛥𝑆𝑎𝑔𝑔

(15)

° As shown in Table 3, the 𝛥𝐺𝑎𝑔𝑔 for SDS-HPMC aggregation process is found to be

strongly negative indicating an energetically favorable process.

° The 𝛥𝐺𝑎𝑔𝑔 for SDS-HPMC

aggregation process is not observed to be sensitive to temperature (Table 3), which may indicate the mechanism of enthalpy-entropy compensation for the SDS-HPMC aggregation process.

Page 25 of 52

° Furthermore, as the 𝛥𝐺𝑎𝑔𝑔 does not provide a complete picture of the free energy for the

SDS-HPMC aggregation process at high SDS concentrations, the phase separation model with the ° ratio of [CAC/CMC] is utilized for determining the free energy (𝛥𝐺𝑇𝑟 ) associated with transfer

of surfactant molecule from micelles to a binding site on the polymer. This equation provides information on the strength of interactions between HPMC and surfactant at a specified temperature (48) and is given below ° ° ° 𝛥𝐺𝑇𝑟 = 𝛥𝐺𝑎𝑔𝑔 − 𝛥𝐺𝑚𝑖𝑐 = 𝑅𝑇 ln

𝐶𝐴𝐶 𝐶𝑀𝐶

(16)

° where 𝛥𝐺𝑇𝑟 is the standard free energy of transfer of one mole of SDS from pure to SDS-

HPMC mixed micelles. For the SDS-HPMC aggregates the ratio was found to be dependent on the polymer concentration and became more negative with increasing polymer concentration. The ° 𝛥𝐺𝑇𝑟 increased from -42.2 ± 0.2 kJ/mol to -43.1 ± 0.2 kJ/mol when the concentration of HPMC

increased from 0.25% to 0.5% w/w. At a higher HPMC concentration, the CAC decreased to a lower SDS concentration while the endothermic maxima increased (Fig. 4). A temperature ° dependence on 𝛥𝐺𝑇𝑟 is observed wherein at 25°C a value of -42.2 kJ/mol is determined that

increases to -45.4 ± 0.1 kJ/mol at 40°C, this further corroborates the predominantly hydrophobic nature of these interactions in this region with the formation of SDS-HPMC aggregates being more ° energetically favored (i.e., more negative than 𝛥𝐺𝑎𝑔𝑔 ). The interaction behavior between SDS ° ° and HPMC consists of a few other thermodynamic components (i.e., 𝐻𝑎𝑔𝑔 , 𝛥𝑆𝑎𝑔𝑔 and 𝛥𝐶𝑝 ) that

are of considerable importance towards understanding mechanism of the interaction process. ° Additionally, the standard enthalpy of aggregation (𝐻𝑎𝑔𝑔 ) is directly obtained from enthalpogram

(𝛥𝐻𝑐𝑜𝑟𝑟 vs. surfactant concentration) and is defined as the standard enthalpy of aggregation per Page 26 of 52

° mole of surfactant monomer unit (Fig. 4b). For the determination of 𝐻𝑎𝑔𝑔 , lines are drawn to fit

the start of interaction or CAC (shown as dotted line A) and the peak point (shown as dotted line ° B) of the endothermic curve, a line is drawn perpendicular to the x axis and 𝛥𝐻𝑎𝑔𝑔 is measured ° from the length of the segment connecting the lines. A similar approach to determine 𝐻𝑎𝑔𝑔 was

reported by Torn et al. (19) while investigating the aggregation behavior between PVP and SDBS ° by calorimetry. 𝛥𝐻𝑎𝑔𝑔 for SDS-HPMC aggregates in this study is determined to be 2.8 kJ/mol

reflecting cooperative interactions between SDS and HPMC that decreased with increasing temperature (Table 3). Heat capacity changes are known to be quantitatively related to conformational changes during the transfer of molecules from an aqueous phase to a non-polar phase (23, 48). 𝛥𝐶𝑝 for the ° HPMC-surfactant interaction process was obtained from the slope of the plot of 𝛥𝐻𝑎𝑔𝑔 vs 𝑇 in

order to evaluate any conformational changes that occur during the aggregation process. 𝛥𝐶𝑝 was determined to be -0.13 kJ/mol K and the negative 𝛥𝐶𝑝 may reflect the transfer of the SDS alkyl chains as well as the non-polar surfaces of HPMC into the more hydrophobic regions of these HPMC-SDS aggregates which is characteristic of hydrophobic interactions (Table 3). This may further suggest that the enthalpic cost of transfer of SDS monomers from aqueous environment to a more hydrophobic environment (SDS-HPMC aggregate) decreases with an increase in the temperature. As discussed in the previous section since the SDS-HPMC aggregation process is not favored enthalpically, an increase of overall entropy is required to compensate unfavorable ° enthalpy of the aggregate formation (56). The standard entropy (𝛥𝑆𝑎𝑔𝑔 ) can be obtained from

equation (13). The breaking of H-bonding network of the water structure at it reorganizes during

Page 27 of 52

° the interaction process may reflect the increase in the overall 𝛥𝑆𝑎𝑔𝑔 for the SDS-HPMC mixed ° system. At 25°C the 𝛥𝑆𝑎𝑔𝑔 is 100.1 J/K mol that decreases to 95.6 J/K mol at 40°C possibly

reflecting smaller conformational (entropic) changes during the cooperative binding of SDS onto HPMC at higher temperatures (Table. 3). The relative gain in entropy of the system as a result of the hydrophobic interactions at higher temperatures is less as the water molecules already possess ° a higher state of disorder(50, 56). Thus, the effect of temperature on 𝛥𝑆𝑎𝑔𝑔 is in agreement with

the hydrophobic effect as a significant driving force for SDS and HPMC interactions. The influence of surfactant headgroup on the surfactant-HPMC interactions is studied by determining the thermodynamic parameters for DTAB-HPMC interactions since the chain length of DTAB is the same as that of SDS. Similar to SDS, the studies were conducted by titrating a micellar solution of DTAB (10x CMC; 129.7mM) into HPMC K-4M solution. The DTAB-HPMC enthalpogram does not show a distinct endothermic peak as seen with the SDS-HPMC (Fig. 6), which is attributed to a lack of or weak interactions between DTAB and HPMC. Weak or a lack of interactions between cationic surfactants and nonionic polymers (i.e., PEO, PVP) have been reported and has been suggested that the larger size of the cation could deter the interaction with nonionic polymers (57). Effect of Molecular Weight of HPMC on the Energetics of HPMC-Surfactant Aggregation Process Molecular weight (MW) of a polymer may influence the energetics of polymer-surfactant interactions that are associated with the conformational changes of polymer to attain most stable aggregate structures (58). In the previous section, it was determined that Csat was proportional to

Page 28 of 52

the concentration of HPMC, which could be attributed to the number of available sites on the HPMC surface for SDS-HPMC interactions to occur. Hence it is reasonable to expect an influence of MW of HPMC on SDS-HPMC interactions. The effect of HPMC MW on the energetics of HPMC-SDS aggregation is presented in Fig. 7 (0.25% w/w HPMC). From the enthalpogram, it is evident that the titration curves for all three MWs of HPMC K-4M, K-15M and K-100M follow the same profile and showed similar shaped curves however, the enthalpogram of the highest MW HPMC (i.e., K-100M) was different from the lower MW HPMC (i.e., K-4M and K-15M). Similar to HPMC K-4M (previous section), the HPMC-SDS enthalpograms for the two higher MW grades of HPMC (K-15M and K-100M) show an endothermic maximum (A→B) followed by an exothermic minimum (B→C) before increasing again and merging with the pure SDS micelles (Cm) dilution curve at approximately zero enthalpy change (C→D) (Fig. 7). The number average MWs (Mn) for HPMC K-4M, K-15M and K-100M are 86 kDa, 120 kDa and >240 kDa, respectively. Moreover, the degree of substitution for all the K-grades of HPMC used in this study consists of the same methyl (19-24%) and hydroxypropyl (7-12%) groups and hence the hydrophobicity of these three HPMC grades is expected to be similar. Since the CAC is known to be sensitive to the hydrophobicity of polymer, it should not be altered with change in HPMC MW (55). As shown in Fig. 7, the CAC for the two lower MW HPMC (K-4M and K-15M) is similar, while the CAC for the higher MW HPMC (K-100M) is slightly lower. The slight decrease in the CAC at higher MW may be attributed to conformational differences between the lower MW and high MW grades of HPMC with the HPMC K-100M having less accessible non-polar surfaces (18). In terms of the other critical concentration parameters, the Csat and Cm of the higher MW HPMC-100M are lower than those of the lower MW HPMC (K-4M and K-15M) Page 29 of 52

(Fig. 7). The Csat and Cm for HPMC K-100M are ~13 mM and 15 mM, respectively as compared to the same for HPMC K-4M are ~17 mM and ~20 mM, respectively and for HPMC K-15M are ~17 mM and ~20 mM, respectively. The HPMC molecular weight is determined to strongly influence the Csat and Cm parameters. To the best of our knowledge, these are the first reported values to show the influence of HPMC MW on the HPMC-SDS aggregate formation and addresses one of the significant gaps (12) in developing an in-depth understanding for this system as identified in the review of the literature in the introduction section. When normalized to HPMC MW, the values of Csat and Cm for HPMC-SDS systems remain distinctly lower for the highest MW HPMC K100M, which indicates that the number of bound SDS is not proportional to the chain length of HPMC. The lower values of Csat and Cm for the highest MW HPMC K-100M suggests that the number of binding sites available for SDS adsorption on HPMC K-100M could be lower than those for the other two lower MW HPMC. This may be attributed to the conformation difference for HPMC K-100M consisting of more buried chains thus providing considerably less binding sites for SDS (16). Dai et. al (16) while investigating the interactions between PEG and SDS also observed a strong influence of MW of PEG on critical concentration parameters. In the PEG-SDS aggregation process an inverse dependence of molecular weight of PEG on the SDS-PEG interaction was attributed to the lower number of available sites for higher MW PEG (55). This hypothesis could be further investigated as potential next steps with the similar ITC experiments using HPMC with increasing hydrophocity (i.e., higher percentage of hydrophobic substituents) where more buried chains could be expected to form during experimentation and be measured by a further decrease in Csat value. In addition binding capacity

Page 30 of 52

and affinity between various HPMC (MW and different substitution) and the surfactants may also prove to be useful in exploring this phenomenon. Influence of Ionic Strength on the Energetics of HPMC-Surfactant Aggregation Process Fig. 8 shows the HPMC-SDS enthalpograms for three ionic strengths at 25°C that were achieved by the addition of 0.1, 0.3 and 0.6% w/w NaCl. Although the amount of NaCl added did not change the shape of the curves significantly, the CAC values decreased significantly to lower SDS concentrations as NaCl concentration increased, which may reflect a decrease in the repulsive forces between SDS molecules promoting the SDS-HPMC aggregation at lower concentrations (35). As a result, the cooperative binding of SDS is enhanced as more SDS molecules are adsorbed onto HPMC. Similarities can be drawn to the results obtained from studying the effect of NaCl on the micellization behavior of two ionic surfactants SDS and DTAB (in the absence of HPMC) in the previous section, where an attractive interaction is observed at low surfactant concentrations with a significant decrease in the CMC of SDS and DTAB at increasing NaCl concentrations. Similar results have also been reported for ionic surfactant micellization behavior in the presence of NaCl where the significant shifts in CMC to lower surfactant values was attributed to the charge ° shielding effect (35, 40, 59). The concentration of NaCl also influences 𝛥𝐻𝑎𝑔𝑔 , as the NaCl ° concentration increases from 0.1% to 0.6% w/w, the values of 𝛥𝐻𝑎𝑔𝑔 increases from 2.1 kJ/mol

to 2.9 kJ/mol at 25°C, respectively, which may suggest that the interactions between HPMC and SDS are stronger at higher NaCl concentrations. This may be attributed to the charge shielding effect. The effect of temperature on the enthalpy of aggregation of SDS-HPMC in the presence of NaCl was also determined and similar to the results presented in the previous section in the absence of NaCl, the endothermic peak decreases with increasing temperature reflecting that the Page 31 of 52

hydrophobic effect could still be the underlying driving force for SDS-HPMC interactions in the presence of NaCl. Furthermore, the influence of NaCl on the exothermic region (B→C) was also evaluated (Fig. 8). The value of Csat was lower at the highest concentration of NaCl (0.6% w/w) may be suggesting a change in conformation of HPMC resulting in a lower number of available binding sites at the highest NaCl concentration. CONCLUSIONS ITC was utilized to study interaction between ionic surfactants and HPMC, SDS-HPMC interactions were determined to be stronger as compared to much weaker interactions between DTAB and HPMC. The driving force for the SDS-HPMC interactions is suggested to be due to the hydrophobic effect, similar to the micellization of SDS. The interaction between SDS and HPMC initiated at a critical concentration of SDS denoted as CAC. From the corrected enthalpy change it was observed that the nature of interaction between SDS and HPMC was an endothermic process that was cooperative in nature and dependent on temperature and ionic strength of the solution.

The negative heat capacity of aggregation for SDS-HPMC along with the effect of

temperature, HPMC molecular weight and ionic strength was utilized to postulate the mechanism of SDS-HPMC aggregate formation. A significant shift in the interaction parameters between HPMC and SDS was observed at highest molecular weight of HPMC (HPMC K-100M) although the general shape of HPMC-SDS interaction curves remained the same for the lower molecular weight grades of HPMC (K-15M and K-4M). At highest molecular weight and concentration of polymer, the critical concentration parameters Csat and Cm are significantly altered and shift to higher concentration of SDS. Ionic strength significantly influenced the SDS-HPMC aggregation process similar to its influence on micellization where the critical concentration parameters (CAC Page 32 of 52

and CMC) decreased with increasing ionic strength for both anionic and cationic surfactant-HPMC systems suggesting stronger interactions. Overall, the interpretation of the microcalorimetric studies at different temperatures and ionic strengths while varying properties of polymer and surfactant was very effective in developing insights into the nature and energetics of HPMC and ionic surfactant interactions.

Acknowledgement The authors would like to thank Dr. Eric Munson, Patrick DeLuca Endowed Professor at the University of Kentucky and Dr. Brad Anderson, H.B. Kostenbauder Professor at the University of Kentucky for the critical review for this paper as well as for the many useful discussions and suggestions. This paper is dedicated to the memory of Dr. Paul M. Bummer.

Page 33 of 52

REFERENCES 1.

Lipinski CA. Drug-like properties and the causes of poor solubility and poor permeability.

Journal of Pharmacological and Toxicological Methods. 2000;44(1):235-49. 2.

Merisko-Liversidge E, Liversidge GG. Nanosizing for oral and parenteral drug delivery:

A perspective on formulating poorly-water soluble compounds using wet media milling technology. Advanced Drug Delivery Reviews. 2011;63(6):427-40. 3.

Murdande SB, Shah DA, Dave RH. Impact of Nanosizing on Solubility and Dissolution

Rate of Poorly Soluble Pharmaceuticals. Journal of Pharmaceutical Sciences. 2015;104(6):2094102. 4.

Elsayed I, Abdelbary AA, Elshafeey AH. Nanosizing of a poorly soluble drug: technique

optimization, factorial analysis, and pharmacokinetic study in healthy human volunteers. International Journal of Nanomedicine. 2014;9:2943-53. 5.

Van Eerdenbrugh B, Van den Mooter G, Augustijns P. Top-down production of drug

nanocrystals: Nanosuspension stabilization, miniaturization and transformation into solid products. International Journal of Pharmaceutics. 2008;364(1):64-75. 6.

Van Eerdenbrugh B, Vermant J, Martens JA, Froyen L, Van Humbeeck J, Augustijns P, et

al. A Screening Study of Surface Stabilization during the Production of Drug Nanocrystals. Journal of Pharmaceutical Sciences. 2009;98(6):2091-103. 7.

Dizaj SM, Vazifehasl Z, Salatin S, Adibkia K, Javadzadeh Y. Nanosizing of drugs: Effect

on dissolution rate. Research in pharmaceutical sciences. 2015;10(2):95-108. 8.

Goddard ED. Polymer Surfactant Interaction.1. Uncharged Water-soluble Polymers and

Charged Surfactants. Colloids and Surfaces. 1986;19(2-3):255-300. Page 34 of 52

9.

Persson K, Wang G, Olofsson G. Self diffussion, thermal effects and viscosity of a

monodisperse associative polymer- self association and interaction with surfactants Journal of the Chemical Society-Faraday Transactions. 1994;90(23):3555-62. 10.

Evertsson H, Nilsson S. Microviscosity in clusters of ethyl hydroxyethyl cellulose and

sodium dodecyl sulfate formed in dilute aqueous solutions as determined with fluorescence probe techniques. Macromolecules. 1997;30(8):2377-85. 11.

Lestari MLAD, Müller RH, Möschwitzer JP. Systematic Screening of Different Surface

Modifiers for the Production of Physically Stable Nanosuspensions. Journal of Pharmaceutical Sciences. 2015;104(3):1128-40. 12.

Singh SK, Nilsson S. Thermodynamics of interaction between some cellulose ethers and

SDS by titration microcalorimetry - I. EHEC and HPMC. Journal of Colloid and Interface Science. 1999;213(1):133-51. 13.

Olofsson G, Wang G. Interactions between surfactants and uncharged polymers in

aqueous-solution studied by microcalorimetry Pure and Applied Chemistry. 1994;66(3):527-32. 14.

Wang YL, Han BX, Yan H, Cooke DJ, Lu JR, Thomas RK. Interaction between

poly(ethylene oxide) and dodecyl sulfates with different monovalent metal counterions studied by microcalorimetry. Langmuir. 1998;14(21):6054-8. 15.

Wang G, Olofsson G. Titration calorimetric study of the interaction between ionic

surfactants and uncharged polymers in aqueous solution. Journal of Physical Chemistry B. 1998;102(46):9276-83.

Page 35 of 52

16.

Dai S, Tam KC. Isothermal titration calorimetry studies of binding interactions between

polyethylene glycol and ionic surfactants. Journal of Physical Chemistry B. 2001;105(44):1075963. 17.

da Silva RC, Loh W, Olofsson G. Calorimetric investigation of temperature effect on the

interaction between poly(ethylene oxide) and sodium dodecylsulfate in water. Thermochimica Acta. 2004;417(2):295-300. 18.

Dai S, Tam KC. Isothermal titration calorimetric studies on the interaction between sodium

dodecyl sulfate and polyethylene glycols of different molecular weights and chain architectures. Colloids and Surfaces A: Physicochemical and Engineering Aspects. 2006;289(1–3):200-6. 19.

Torn LH, de Keizer A, Koopal LK, Lyklema J. Titration microcalorimetry of

poly(vinylpyrrolidone) and sodium dodecylbenzenesulphonate in aqueous solutions. Colloids and Surfaces a-Physicochemical and Engineering Aspects. 1999;160(3):237-46. 20.

Torn LH, de Keizer A, Koopal LK, Lyklema J. Mixed adsorption of poly(vinylpyrrolidone)

and sodium dodecylbenzenesulfonate on kaolinite. Journal of Colloid and Interface Science. 2003;260(1):1-8. 21.

Blandamer MJ, Cullis PM, Gleeson PT. Three important calorimetric applications of a

classic thermodynamic equation. Chemical Society Reviews. 2003;32(5):264-7. 22.

Blandamer MJ, Cullis PM, Engberts J. Titration microcalorimetry. Journal of the Chemical

Society-Faraday Transactions. 1998;94(16):2261-7. 23.

Freyer MW, Lewis EA. Isothermal titration calorimetry: Experimental design, data

analysis, and probing Macromolecule/Ligand binding and kinetic interactions. In: Correia JJ,

Page 36 of 52

Detrich HW, editors. Biophysical Tools for Biologists: Vol 1 in Vitro Techniques. Methods in Cell Biology. 842008. p. 79-113. 24.

Yoshida E. Self-assembly of poly(allylamine hydrochloride) through electrostatic

interaction with sodium dodecyl sulfate. Colloid and Polymer Science. 2010;288(12-13):1321-5. 25.

van Stam J, Depaemelaere S, De Schryver FC. Micellar aggregation numbers - A

fluorescence study. Journal of Chemical Education. 1998;75(1):93-8. 26.

Sovilj V, Petrovic L. Interaction and phase separation in the system HPMC/NaCMC/SDS.

Colloids and Surfaces a-Physicochemical and Engineering Aspects. 2007;298(1-2):94-8. 27.

Woolley EM, Burchfield TE. Model for Thermodynamics of Ionic Surfactant Solutions.2.

Enthalpies, Heat -capacities and Volumes Journal of Physical Chemistry. 1984;88(10):2155-63. 28.

Blandamer MJ, Briggs B, Cullis PM, Engberts J. Titration microcalorimetry of mixed

alkyltrimethylammonium bromide surfactant aqueous solutions. Physical Chemistry Chemical Physics. 2000;2(22):5146-53. 29.

Beyer K, Leine D, Blume A. The demicellization of alkyltrimethylammonium bromides in

0.1 M sodium chloride solution studied by isothermal titration calorimetry. Colloids and Surfaces B-Biointerfaces. 2006;49(1):31-9. 30.

Chatterjee A, Moulik SP, Sanyal SK, Mishra BK, Puri PM. Thermodynamics of Micelle

Formation of Ionic Surfactants:  A Critical Assessment for Sodium Dodecyl Sulfate, Cetyl Pyridinium Chloride and Dioctyl Sulfosuccinate (Na Salt) by Microcalorimetric, Conductometric, and Tensiometric Measurements. The Journal of Physical Chemistry B. 2001;105(51):12823-31.

Page 37 of 52

31.

Paula S, Sus W, Tuchtenhagen J, Blume A. Thermodynamics of micelle formation as a

function of temperature - A high sensitivity titration calorimetry study Journal of Physical Chemistry. 1995;99(30):11742-51. 32.

Hait SK, Majhi PR, Blume A, Moulik SP. A critical assessment of micellization of sodium

dodecyl benzene sulfonate (SDBS) and its interaction with poly(vinyl pyrrolidone) and hydrophobically modified polymers, JR 400 and LM 200. Journal of Physical Chemistry B. 2003;107(15):3650-8. 33.

Horin S, Arai H. Effect of added salt on interaction between polymer and detergent in

aqueous solution. Journal of Colloid and Interface Science. 1970;32(3):547-&. 34.

Hu YZ, Zhao CL, Winnik MA, Sundararajan PR. Fluorescence studies of the interaction

of sodium dodecyl-sulfate with hydrophobically modified poly(ethylene oxide). Langmuir. 1990;6(4):880-3. 35.

Phillips JN, Mysels KJ. Light scattering by aqueous solutions of sodium lauryl sulfate.

Journal of Physical Chemistry. 1955;59(4):325-30. 36.

Blandamer MJ, Cullis PM, Soldi LG, Rao KC, Subha MCS. Effects of added DOTAB on

the cmc and enthalpy of micelle formation at 298.2 K for CTAB(aq). Journal of Thermal Analysis. 1996;46(6):1583-8. 37.

Binanalimbele W, Zana R. Electrical-conductivity study of the self association of ionic

surfactants in solution in ethyleneglycol, formic acid and formamide Colloid and Polymer Science. 1989;267(5):440-7.

Page 38 of 52

38.

Thomas DC, Christian SD. Micellar and surface behavior of sodium deoxycholate

characterized by surface-tension and ellipsometry methods Journal of Colloid and Interface Science. 1980;78(2):466-78. 39.

Martin AN, Bustamante P. Physical Pharmacy: Physical Chemical Principles in the

Pharmaceutical Sciences: Lea & Febiger; 1993. 40.

Volpe PLO, Silva EA. Calorimetric study of SDS micelle formation in water and in NaCl

solution at 298-K. Thermochimica Acta. 1995;257:59-66. 41.

Johnson I, Olofsson G. Solubilization of 2-Propanol and Pentanol in sodium dodecyl-

sulfate micelles- A thermochemical study Journal of Colloid and Interface Science. 1987;115(1):56-64. 42.

Kroflic A, Sarac B, Bester-Rogac M. Influence of the alkyl chain length, temperature, and

added salt on the thermodynamics of micellization: Alkyltrimethylammonium chlorides in NaCl aqueous solutions. Journal of Chemical Thermodynamics. 2011;43(10):1557-63. 43.

Kessler A, Zeeb B, Kranz B, Menendez-Aguirre O, Fischer L, Hinrichs J, et al. Isothermal

titration calorimetry as a tool to determine the thermodynamics of demicellization processes. Review of Scientific Instruments. 2012;83(10). 44.

Blandamer MJ, Briggs B, Cullis PM, Irlam KD, Engberts J, Kevelam J. Titration

microcalorimetry of adsorption processes in aqueous systems - Interaction of sodium dodecylsulfate and sodium decylsulfate with poly(N-vinylpyrrolidone). Journal of the Chemical Society-Faraday Transactions. 1998;94(2):259-66.

Page 39 of 52

45.

Sarac

B,

Bester-Rogac

M.

Temperature

and

salt-induced

micellization

of

dodecyltrimethylammonium chloride in aqueous solution: A thermodynamic study. Journal of Colloid and Interface Science. 2009;338(1):216-21. 46.

Bouchemal K, Agnely F, Koffi A, Djabourov M, Ponchel G. What can isothermal titration

microcalorimetry experiments tell us about the self-organization of surfactants into micelles? Journal of Molecular Recognition. 2010;23(4):335-42. 47.

Sovilj

VJ,

Petrovic

LB.

Influence

of

hydroxypropylmethyl

cellulose-sodium

dodecylsulfate interaction on the solution conductivity and viscosity and emulsion stability. Carbohydrate Polymers. 2006;64(1):41-9. 48.

La Mesa C. Polymer-surfactant and protein-surfactant interactions. Journal of Colloid and

Interface Science. 2005;286(1):148-57. 49.

Nilsson S. Interactions between Water-Soluble Cellulose Derivatives and Surfactants. 1.

The HPMC/SDS/Water System. Macromolecules. 1995;28(23):7837-44. 50.

Dill KA, Privalov PL, Gill SJ, Murphy KP. The Meaning of Hydrophobicity. Science.

1990;250(4978):297-8. 51.

.

52.

Abraham MH. Thermodynamics of solution of homologous series of solutes in water

Journal of the Chemical Society-Faraday Transactions I. 1984;80:153-81. 53.

Nichols N, Skold R, Spink C, Suurkuusk J, Wadso I. Additivity relations for heat-capacities

of non-electrolytes in aqueous-solution Journal of Chemical Thermodynamics. 1976;8(11):108193.

Page 40 of 52

54.

Johnson I, Olofsson G, Landgren M, Jonsson B. Solubilization of Pentanol in sodium

dodecyl-sulfate micelles- Interpretation of calorimetric results using a theoretical model Journal of the Chemical Society-Faraday Transactions I. 1989;85:4211-25. 55.

Dai S, Tam KC. Isothermal titration calorimetric studies on the temperature dependence of

binding interactions between poly(propylene glycol)s and sodium dodecyl sulfate. Langmuir. 2004;20(6):2177-83. 56.

Rekharsky MV, Inoue Y. Complexation thermodynamics of cyclodextrins. Chemical

Reviews. 1998;98(5):1875-917. 57.

Benrraou M, Bales B, Zana R. Effect of the nature of the counterion on the interaction

between cesium and tetraalkylammonium dodecylsulfates and poly(ethylene oxide) or poly(vinylpyrolidone). Journal of Colloid and Interface Science. 2003;267(2):519-23. 58.

Evertsson H, Nilsson S, Holmberg C, Sundelof LO. Temperature effects on the interactions

between EHEC and SDS in dilute aqueous solutions. Steady-state fluorescence quenching and equilibrium dialysis investigations. Langmuir. 1996;12(24):5781-9. 59.

Cabane B. Structure of some polymer-detergent aggregates in water. The Journal of

Physical Chemistry. 1977;81(17):1639-45. 60.

Umlong IM, Ismail K. Micellization behaviour of sodium dodecyl sulfate in different

electrolyte media. Colloids and Surfaces A: Physicochemical and Engineering Aspects. 2007;299(1–3):8-14.

Page 41 of 52

FIGURES

Fig.1. Schematic of ITC setup for the power compensation mechanism and raw data for a general case (23).

Page 42 of 52

C

ΔHmic

Fig. 2. Representative data transformation to determine critical concentrations: (a) raw data directly obtained from isothermal titration calorimetry (heat Flow (µW) vs. time), (b) enthalpograms depicting enthalpy (ΔHapp) change as a function of total surfactant (i.e., CTAB) concentration, and (c) the first derivative of the curve (b).

Page 43 of 52

( a)

( b)

° Fig. 3. Microcalorimetric determination of 𝚫𝐇𝐦𝐢𝐜 of (a) SDS and (b) DTAB at 25°C (circles), 32°C (triangles), 40°C (diamonds) and 50°C (squares). The titration at each temperature consisted of 25µL aliquots of surfactant solution (SDS~ 87.38 mM and DTAB~ 129.72 mM) into 2.5 ml water in 50 steps.

Page 44 of 52

Fig. 4. Plot of (a) apparent enthalpy (ΔHapp) and (b) corrected enthalpy (ΔHcorr) as a function of total SDS concentration in the presence of (squares) 0% HPMC K-4M, (diamonds) 0.25% HPMC K-4M and (triangles) 0.5% HPMC K-4M at 25°C.

Page 45 of 52

Fig. 5. Effect of temperature on apparent enthalpy (ΔHapp) as a function of total SDS concentration in the presence of HPMC K-4M at 25°C (squares; 0.25% HPMC), 32°C (triangles; 0.25% HPMC) and 40°C (crosses; 0.25% HPMC).

Page 46 of 52

Fig. 6. Plot of apparent enthalpy (ΔHapp) as a function of total DTAB concentration in the presence of (squares) 0% HPMC K-4M and (diamonds) 0.25% HPMC K-4M at 25°C.

Page 47 of 52

Fig. 7. Plot of apparent enthalpy (ΔHapp) as a function various molecular weight of HPMC at 25°C. 0.25% HPMC K-4M, (triangles) 0.25% HPMC K-4M, (squares) 0.25% HPMC K15 and (circles) 0.25% HPMC K-100.

Page 48 of 52

Fig. 8. Enthalpogram of the effect of NaCl on aggregation behavior at 25°C for 0.25% HPMC/SDS at 0 % NaCl, (cross), 0.1% NaCl (squares), 0.3% NaCl (triangles) and 0.6% NaCl (diamonds).

Page 49 of 52

TABLES Table 1: Thermodynamic parameters for SDS micellization

° 𝛥𝐻𝑚𝑖𝑐

° 𝛥𝐻𝑚𝑖𝑐

° 𝛥𝐺𝑚𝑖𝑐

𝛥𝐶 p

° 𝛥𝑆𝑚𝑖𝑐

Temp. CMC

CMC

⁰C

(mM)

(mM)

kJ/mol

kJ/mol

kJ/mol

J/°K

25⁰ C

8.3

7.2-8.4a

-0.6

-0.5a

-23.7

77.1

kJ/°mol K

8.1b(8.1c) 32⁰ C

8.2(±0.2)

--

-3.5(±0.1)

--

-24.3(±0.2)

67.7

40⁰ C

7.7(±0.2)

7.9-8.7a

-7.5(±0.0)

-7.1a

-25.3(±0.2)

56.2

50⁰ C

8.1(±0.0)

8.1-9.2a

-10.7(±0.3)

-11.7a

-25.8(±0.0)

45.9

a

ITC generated values; Ref (12)

b

values of cmc from surface tension method; Ref (60)

c

values of cmc from conductance method are shown in the parentheses; Ref (60)

Page 50 of 52

-0.5

Table 2: Thermodynamic parameters for DTAB micellization Temp.

CMC

CMCa

° 𝛥𝐻𝑚𝑖𝑐

⁰C

(mM)

(mM)

kJ/mol

kJ/mol

kJ/mol

J/°K

kJ/°mol K

25⁰ C

14.5(±0.4)

13.5

1.9a

-1.7(±0.1)

-20.9(±0.1)

64.3(±0)

-0.4

32⁰ C

14.9(±0.3)

--

--

-4.4(±0.0)

-21.3(±0.1)

53.2(±0)

40⁰ C

15.9(±0.3)

--

--

-7.4(±0.1)

-21.5(±0.3)

41(±0)

a

° 𝛥𝐻𝑚𝑖𝑐

Ref (29)

Page 51 of 52

° 𝛥𝐺𝑚𝑖𝑐

° 𝛥𝑆𝑚𝑖𝑐

𝛥𝐶p

Table 3: Thermodynamic parameters for SDS-HPMC aggregation HPMC Conc.

Temp.

CACa

% w/w

⁰C

mM

kJ/mol

kJ/mol

J/K

25⁰ C

4.1(±0.2)

2.4 (±0.1)

-27.1(±0.3)

100.1

32⁰ C

3.5b(±0.1)

0.9 (±0.2)

-28.6(±0.1)

97.1

40⁰ C

3.4b(±0.4)

0.4(±0.2)

-29.5(±0.6)

95.6

25⁰ C

4.2(±0.2)

2.5(±0.1)

-27.1(±0.2)

99.5

0.25 0.5 a

° 𝛥𝐻𝑎𝑔𝑔

° 𝛥𝐺𝑎𝑔𝑔

CAC: critical aggregation concentration

b

approximated as presence of distinct CAC is absent

Page 52 of 52

° 𝛥𝑆𝑎𝑔𝑔

𝛥𝐶p kJ/°K mol -0.13