Thermodynamic Studies of Ionic Hydration and Interactions for Amino

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Thermodynamic Studies of Ionic Hydration and Interactions for Amino Acid Ionic Liquids in Aqueous Solutions at 298.15 K Dilip H Dagade, Kavita R. Madkar, Sandeep P. Shinde, and Seema S. Barge J. Phys. Chem. B, Just Accepted Manuscript • Publication Date (Web): 07 Jan 2013 Downloaded from http://pubs.acs.org on January 10, 2013

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The Journal of Physical Chemistry

Thermodynamic Studies of Ionic Hydration and Interactions for Amino Acid Ionic Liquids in Aqueous Solutions at 298.15 K Dilip H. Dagade,* Kavita R. Madkar, Sandeep P. Shinde, Seema S. Barge Department of Chemistry, Shivaji University, Kolhapur – 416004, India.

ABSTRACT: Amino acid ionic liquids are special class of ionic liquids due to their unique acidbase behaviour, biological significance and applications in different fields such as templates in synthetic chemistry, stabilizers for biological macromolecules, etc. The physicochemical properties of these ionic liquids can easily be altered by making the different combinations of amino acids as anion along-with possible cation modification which makes amino acid ionic liquids more suitable to understand the different kinds of molecular and ionic interactions with sufficient depth so that they can provide fruitful information for molecular level understanding of more complicated biological processes. In this context, volumetric and osmotic coefficient measurements for aqueous solutions containing 1-ethyl-3-methylimidazolium ([Emim]) based amino acid ionic liquids of glycine, alanine, valine, leucine and isoleucine are reported at 298.15 K. From experimental osmotic coefficient data, mean molal activity coefficients of ionic liquids were estimated and analysed using Debye-Hückel and Pitzer models. The hydration numbers of ionic liquids in aqueous solutions were obtained using activity data. Pitzer ion interaction parameters are estimated and compared with other electrolytes reported in literature. The nonelectrolyte contribution to the aqueous solutions containing ionic liquids were studied by calculating osmotic second virial coefficient through an application of McMillan-Mayer theory 1

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of solution. It has been found the second osmotic virial coefficient which includes volume effects correlates linearly with the Pitzer ion interaction parameter estimated independently from osmotic data as well as the hydrophobicity of ionic liquids. Enthalpy-entropy compensation effect, explained using the Starikov-Nordén model of enthalpy-entropy compensation, and partial molar entropy analysis for aqueous [Emim][Gly] solutions are made by using experimental Gibb’s free energy data and literature enthalpy data. This study highlights that the hydrophobic interaction persists even in the limit of infinite dilution where the hydration effects are usually dominant implying importance of hydrophobic hydration. Analysis of the results further shows that the hydration of amino acid ionic liquids occurs through the cooperative H-bond formation with kosmotropic effect in contrast to the usual inorganic salts or hydrophobic salts like tetraalkylammonium halides. Keywords: Osmotic coefficient, hydration number, hydrophobic hydration, enthalpy-entropy compensation, kosmotropic effect, cooperative H-bonding, hydrophobicity. 1. INTRODUCTION The room temperature molten salts generally called as room temperature ionic liquids (RTILs) have attracted researchers from various fields due to unique physicochemical properties of these RTILs such as size, shape, polarity, hydrophobicity, solvent miscibility, etc. which can be tuned as per the need by making the possible combinations of cations and anions (estimate shows ~ 106 combinations of known cations and anions can form ionic liquids).1-8 This has led to the development of wide variety of ionic liquids (ILs) having their applications in the field of nanoscience, electrochemical sensors, supercapacitors, material science, biocatalytic reactions, biosensors, biopreservation as well as protein solubilization, stabilization and crystallization.9-15 2


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One of the important class of ionic liquids called as amino acid ionic liquids (AAILs) has emerged due to classic work of Fukumoto et al.16 in which authors have synthesized imidazolium based AAILs from 20 natural amino acids. This work simulated the field further resulting into synthesis of series of amino acid ionic liquids with different combinations of cations and anions (amino acid moiety can be used as either cation or anion) which introduced the concept of bioionic liquids in the field of ILs where both the cation and anion are of biological origin.17-23 Use of amino acid ionic liquids or bio-ionic liquids instead of traditional ionic liquids is more demanding due to their biodegradability and nontoxicity which can play an important role in pharmaceuticals for drug formulations as well as in protein chemistry to understand the natural osmolyte effects, protein folding, etc. and as an intermediates or templates in synthesis of chiral compounds, peptides, functional materials, etc. However, to explore and understand the potential of these ILs, it is essential to understand the nature of interactions responsible for their existence in pure liquid state at ambient conditions along-with the ionic interactions in aqueous solutions containing these ionic liquids e.g. amongst various non-covalent interactions such as H-bonding, van der Waals forces, ion-pairing, ionic and hydrophobic hydration etc. whether the particular kind of interaction is responsible or their collective and cooperative effect is responsible for overall behaviour. In this context, many researchers have studied the physicochemical properties of amino acid ionic liquids in their pure state24-34 as well as in aqueous solutions.35-44 The native structures and conformations of biological macromolecules like proteins, enzymes etc. responsible for their proper functioning is believed to be dominantly controlled by non-covalent interactions where hydrophobic interactions and H-bonding plays a major role. To understand these interactions at molecular level, the building blocks of such biopolymers i.e. natural amino acids and their aqueous solutions have been well studied experimentally as well as 3



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theoretically.45-57 However, it seems interesting that the limited solubility of these natural amino acids in aqueous and even in non-aqueous media gets increased many-fold when amino acids are converted into ionic liquids. Further, this transformation completely alters the acid-base behaviour of amino acids which may affect the solution properties. If such a medium is used for dissolution and stabilization of biomacromolecules, it will be very interesting to see how the individual amino acid units from ionic liquids interact with the amino acids of proteins or enzymes, carbohydrate units of DNA or RNA etc. since the amino acid moiety in AAIL has ability to form strong H-bonds26 which can help to increase the solubility and stability of biomolecules containing carbohydrates particularly glycoproteins, DNA, RNA etc. Further, the molecular recognition of amino acids when they are present in proteins as a backbone species or they are in the side chains will be of great significance in protein chemistry and probably the amino acid ionic liquids can play an important role in this direction. However, due to competitive hydration of proteins and osmolytes or denaturants, the understanding of role of ionic liquids at molecular level as a molecular recogniser, denaturant, osmolyte etc. is difficult when there is lack of detailed thermodynamic understanding of aqueous solutions amino acid ionic liquids. In this context and keeping in mind the potential use of AAILs as mentioned above, we are reporting here the detailed thermodynamic study of some 1-ethyl-3-methylimidazolium based amino acid ionic liquids through experimental volumetric and osmotic pressure measurements. 2. EXPERIMENTAL METHODS Glycine (Alfa Aesar, >99%), L-Alanine (Himedia, >99%), L-Valine (Himedia, >99%), L-leucine (Alfa Aesar, >99%) and L-Isoleucine (Himedia, >99%), 1-methylimidazole (Spectrochem, >99 4


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%), ethyl bromide (Spectrochem, >99 %) and toluene (Merck, >99 %) were used without further purification. The salt NaCl of AR grade (BDH, >99.9%) was dried under vacuum at 393 K for 24 hrs before use. Quartz doubly distilled water was used for entire experimentation. 1-ethyl-3-methylimidazolium bromide was synthesized through quaternization of 1methylimidazole using ethyl bromide for which toluene was used as solvent. The reaction mixture was kept at room temperature for 10 hrs which results into separation of two liquid layers. After the completion of reaction, the solvent toluene was decanted and the product was washed with ethyl acetate to remove any unreacted reactant. Finally, the traces of ethyl acetate was removed under vacuum drying at 60 °C for 12 hrs and then it was further kept under vacuum at 100 °C for 24 hrs to remove any traces of water. The water content in the final product was found to be 0.2 mass % with the help of Karl-Fischer Titration. 1-ethyl-3-methylimidazolium ([Emim]) based Amino acid ionic liquids namely [Emim][Gly], [Emim][Ala], [Emim][Val], [Emim][Leu], [Emim][Ile] were synthesized using the reported method of Fukumoto et al.16 where the 1-ethyl-3-methylimidazolium hydroxide ([Emim][OH]) was prepared by passing aqueous solution of 1-ethyl-3-methylimidazolium bromide ([Emim][Br]) through the –OH form anion exchange resin which was freshly generated from Amberlite IRA-400 Cl using 2% aqueous sodium hydroxide solution. An aqueous [Emim][OH] solution was added drop wise to slightly excess equimolar aqueous amino acid solution and the resultant mixture was stirred under the cooling for 12 hrs. Water was then evaporated by heating the neutralized mixture at 45-50 °C. The excess amino acids were then removed through precipitation by adding the mixture of acetonitrile–methanol (in 9:1 ratio) under vigorous stirring. Finally, the filtrate was evaporated to remove solvent and the product 5



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left was dried under vacuum for 2 days at 80 °C. No precipitation was observed when test was performed with silver nitrate ensuring the absence of any traces of bromide in the final product to the detection level. pH-metric titration was also performed to ensure that no free OH− ions are present in the final product. The water content in the ionic liquids was found to be less than 0.5 mass % and was taken into consideration during preparation of solutions. All the solutions were prepared on molality basis and whenever necessary converted to molarity scale using density data at 298.15 K. A Mettler Toledo AB204-S balance having readability of 0.1 mg was used for weighing. The density measurements were made by using Anton Paar digital densitometer (model: DMA 60/602) at 298.15±0.02 K for which Julabo F-25 MP cryostat was used for temperature control. After applying humidity and lab pressure corrections the uncertainty in the density measurements was found to be ±5×10−3 kg.m-3 The osmotic coefficients () of aqueous AAIL solutions were determined using KNAUER K-7000 vapor pressure osmometer at 298.15±0.001 K. The uncertainty in osmotic coefficient data was found to be ±1×10−3. The details about calibration, measurements and error analysis were reported earlier.53,58,59 3. RESULTS a. Volumetric properties The apparent molar volumes ( ) and partial molar volumes of solute () and solvent water () for aqueous solutions of amino acid ionic liquids (AAILs) at 298.15 K were calculated from experimental density data using the usual equations.60,61 The data of density, , and for

6


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aqueous solutions of amino acid ionic liquids (AAILs) at 298.15 K are reported in Table 1. The data can be fitted with concentration c (in mol.m−3) using the equation √

(1)

where is the limiting apparent molar volume, SV is the Debye-Hückel limiting slope (SV = 1.868 × 106 (mol.mm−3)−3/2 for aqueous solutions of 1:1 electrolytes62 at 298.15 K and BV is a deviation parameter. and BV values are given in Table 2 for studied AAILs which was obtained by plotting ( √ ) against c as shown in Figure 1. b. Osmotic and activity coefficients Experimental osmotic coefficient () data obtained in the concentration range of 0.01 to 0.6 kg.mol−1 for AAILs in aqueous solutions at 298.15 K (see Table 3 and Figure 2) are used to calculate water activities (aw) using the equation

ln

(2)

where x1 and x2 are the mole fractions of solvent and solute, respectively. This data of water activity were used to obtain activity coefficients of water (γ1) in aqueous solution of AAILs. Further, the experimental osmotic coefficient data have been expressed as 1

.

√ ∑!" ⁄

(3)

where Aγ is the Debye-Hückel limiting slope for aqueous solutions (Aγ is 0.5115 at 298.15 K). The coefficients Ai in eqn (3) were obtained by method of least square fit and are given in Table

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4. The mean molal activity coefficients of AAILs (γ±) in aqueous solutions were obtained using equation63 ln #$ 2.303 √ ∑!"

(

⁄

(4)

The activity coefficient data which have been converted to mole fraction scale are used to calculate excess Gibbs free energy change (∆GE) using the following equation ∆* + ,-. ∑ / ln #

(5)

where ν is the number of ions produced on dissociation of electrolyte, R is gas constant, T is absolute temperature and, xi and γi are respectively the mole fraction and activity coefficient of ith component in the mixture. The data of water activity, activity coefficient of components and excess Gibbs free energy change due to mixing are included in Table 3. The concentration dependence of mean molal activity coefficients of AAILs is shown in Figure 3. The hydration numbers (h) for AAILs have been calculated using the mean molal activity coefficient data and following Robinson and Stokes method64 according to which for aqueous 1:1 electrolyte solutions the γ± can be expressed as 2 √4

3 log #$ (.567

√4

8

9 log log:1 0.018?

(6)

Here, a is the distance of the closest approach of oppositely charged ions. The estimated hydration numbers for studied ionic liquids are given in Table 5. c. Enthalpy-entropy compensation and partial molal entropies for aqueous solutions of [Emim][Gly] 8


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Enthalpy-entropy compensation (EEC) is now a physically-chemically valid phenomenon.65-72 Many researchers have applied this to large numbers of systems ranging from small molecules like amino acids, salts, etc. to large biomolecules like proteins, DNA, etc. to understand the molecular interactions in solution phase.65-75 The EEC effect can yield the useful valid data-set only when the independent experimental source of enthalpy and Gibbs free energy data is available.69-72 We earlier reported an analysis of activity and heat content data to understand EEC effects and to obtain partial molar entropies of solute as well as solvent for model compounds like 18-crown-6, 15-crown-5, α-cyclodextrin as well as for amino acids and ionic liquids in aqueous solutions.52,63,73-75 Similar methodology has been adopted here for analysis of activity data of aqueous [Emim][Gly] solutions for which the literature heat data is available in low concentration region. For other systems studied in this work, no heat data is available in literature and hence the analysis is limited to only aqueous [Emim][Gly] solutions at 298.15 K. For aqueous [Emim][Gly] solutions, molar enthalpy of solution (@4 ) data are available in literature76 along-with standard molar enthalpy of solution (@4 ) at 298.15 K. From literature

enthalpy data76, the relative apparent enthalpy of solution (A ) for aqueous [Emim][Gly] at , and further used to obtain 298.15 K were obtained using the expression A @4 @4

enthalpy of mixing (∆Hmix) per mole of solution using the relation77 ∆BCDE

/

(7)

where n1 and n2 are the number of moles of solvent i.e. water (55.51) and of [Emim][Gly], respectively, while L2 is relative partial molar heat content of pure solute, which is the molar solution enthalpy at infinite dilution but with opposite sign (i.e. F @4 ). L2 is taken from the

9



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literature76 as 34.34 kJ.mol−1 at 298.15 K. The data of ∆Hmix thus obtained were combined with ∆Gmix (Δ*4 ν-. ∑ / ln where xi and ai are the mole fraction and activity of ith component in the mixture) data from the present work to obtain T∆Smix values at different concentrations using the relation Δ*4 Δ@4 .Δ 4 . Excess entropy due to mixing is calculated using the excess Gibbs free energy data (Δ* + ) and excess enthalpy (Δ@ + Δ@4 ) data. The data of excess and mixing thermodynamic properties for aqueous [Emim][Gly] solutions are given in Table 6. Concentration dependence of mixing thermodynamic properties is shown in Figure 4 while Figure 5 represents the enthalpy-entropy compensation effect. The relative partial molar heat content of the solvent (FI ) and [Emim][Gly] (FI ) values were determined using the equations77 J⁄

4 LM FI KK.K9 L4N⁄

(8)

O,Q,!

⁄

4 LM FI A 9 L4N⁄

O,Q,!

(9)

The partial molar enthalpies thus calculated have been used along-with activity data to obtain partial molar entropies of solvent water < ̅ > and [Emim][Gly] < ̅ > with the help of following equations77,78 < ̅ >

AI

< ̅ >

TAI UA V

O

- ln S

O

(10)

,- ln

10


(11)

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where is the molar entropy of pure liquid water and is molar entropy of [Emim][Gly] in a hypothetical ideal solution of unit mole fraction. Excess partial molar entropies of water < ̅ >+ and [Emim][Gly] < ̅ >+ were further calculated using the relations77,78 < ̅ >+

AI

< ̅ >+

TAI UA V

O

- ln #

O

,- ln #$

(12)

(13)

The data of apparent and relative partial molal enthalpies, partial molar entropies and excess partial molar entropies of components are also included in Table 6. The concentration dependence of excess partial molar entropies of water and [Emim][Gly] for aqueous solutions is shown in Figure 7. 4. DISCUSSION a. Volumetric properties The apparent molar volume ( ) for the studied AAILs in aqueous solutions at 298.15 K increases slightly with increase in concentration of ionic liquid (see Table 1 and Figure 1) except for [Emim][Gly] system where the effect is more pronounced as observed in case of alkali halides. The increase in with concentration shows that the ionic hydration is dominant as compared to ion-ion association. The sign and magnitude of deviation parameter BV is generally attributed to the water structuring effects for aqueous electrolytic solutions.62 However, the direct correlation of BV with molecular interactions is difficult and hence at this point we do not offer any further explanation for the observed trends of BV values for the studied ionic liquids in aqueous solutions. The excess partial molar volume (I+ ) is estimated using the partial molar 11



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volume of ionic liquids at infinite dilution (I ) in aqueous solutions at 298.15 K and the molar volume of pure ionic liquid (V2) calculated with help of literature density data of amino acid ionic liquids in pure state at 298.15 K. The literature density data at 298.15 K is available for pure [Emim][Gly],25 [Emim][Ala]79 and[Emim][Val]80 and hence the calculations of I+ are limited to only these three AAILs. The estimated I+ values are −9.44 × 103 mm3.mol−1, −8.93 × 103 mm3.mol−1, −3.71 × 103 mm3.mol−1. The negative I+ values indicate that solute molecules occupy less volume in aqueous solution than that in its pure state at the same temperature. Freidman and Scheraga81 and Franks and Smith82 reported negative values of I+ for aqueous alcohol solutions which have explained by Hagler et al83 and Franks84 in terms of hydrophobic hydration which occurs with economy of space. Thus, seen in this light, the dissolution of AAILs in aqueous solutions occurs with hydrophobic hydration. b. Osmotic and Activity Coefficients Examination of Figure 2 shows that the osmotic coefficient data for studied aqueous solutions of amino acid ionic liquids follow Debye-Hückel limiting law in low concentration region i.e. below ~ 0.02 mol.kg−1 whereas at higher concentrations positive deviation from limiting law exists. The osmotic coefficient decreases with concentration in low concentration region where hydration effects are dominant and passes through a minimum around 0.3 mol.kg−1 after which the ion-ion interactions gets dominated over hydration effects. As anion becomes more hydrophobic due to additional alkyl groups, the ion-ion interactions become more favourable due to hydrophobic association of anions leading to increase of osmotic coefficient at higher concentrations with shifting the minimum in towards lower concentration. Further, the increase in size of [AA]− on going from glycine to leucine reduces the charge density of anions 12


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which favours hydrophobic ion-ion association as reflected in concentration dependence of osmotic coefficients for the studied ionic liquids (Figure 2). Similar effects can also be seen from the concentration dependence of mean molal activity coefficients of AAILs (Figure 3). Generally, the increase in size of the hydrophobic cations increases the activity coefficient as observed in tetraalkylamonium fluoride salts whereas the increase in size of anions on going from fluoride to iodide decreases the activity coefficients as shown in case of tetrapropylammonium salts.85 However, in present case the activity coefficient increases with increase in anion size which must be largely due to the increase of hydrophobicity of anions having H-bonding ability resulting into dominance over charge density effects. The mean molal activity coefficient of AAILs have been fitted using the eqn (6) to obtain hydration number h and ion size parameter a. The estimated hydration numbers of AAILs, as reported in Table 5, are larger in magnitude as compared to that of simple monovalent electrolytes like alkali halide or tetraalkylammonium halides. Further, our previous analysis for [Emim][Br] and [Bmim][Cl] ionic liquids in aqueous solutions at 298.15 K showed that the imidazolium cations does not get hydrated63 and hence the observed hydration numbers are due to the anion hydration. Thus, the reported hydration numbers are due to the amino acid anions and values are greater than that of corresponding zwitterionic amino acids in aqueous solutions estimated using different methods.86 Since the imidazolium cations exist in aqueous solutions with peripheral hydration like tetraalkylammonium ions, the number of H-bonds must increase in the secondary hydration shell of cation, i.e. region B in Frank and Wen’s mixture model, which should lead to kosmotropic effect (water structure making effect). Seen in this light, amino acid ionic liquids should behave as structure breakers due to strong anionic hydration. However, this is not observed in the these ionic liquids in aqueous solution due to the ability of amino acid 13



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anions to form cooperative H-bonds with bulk water and persistent of hydrophobic effect even at infinite dilution as evidenced from the detailed entropic analysis of [Emim][Gly] solutions discussed below. The magnitude of ion size parameter a which is very small for the [Emim][Gly] system suggests that even in the low concentration limit, the oppositely charged ions are very closer probably the co-sphere overlap of hydrated ions can occur. c. Application of McMillan-Mayer theory An application of McMillan-Mayer theory of solutions to aqueous electrolytic solutions yields information about relative nonelectrolyte contribution to the solute-solvent and solute-solute interactions.87,88 In dilute solutions, if the Debye-Hückel electrostatic contribution is subtracted from thermodynamic parameter e.g. lnγ±, then the remainder will be linear in molality similar to what is being observed for aqueous nonelectrolyte solutions. Accordingly, the mean molal activity coefficient of the solute (γ±) in the dilute concentration range can be represented as ln #$ W⁄ T1 X⁄ V

U

Y

(14)

where α =1.173 kg1/2.mol−1/2 at 298.15 K, b=1.0 kg1/2.mol−1/2 and ω is the nonelectrolyte solutesolute interaction parameter. According to McMillan and Mayer89 Z

[O

∗ ∗ \ \ \ ⋯

(15)

∗ ∗ where n is the number density of the solute and and are the osmotic second and third

virial coefficient for solute-solute interaction. Total Gibbs free energy of solution can be expressed as a function of mole ratio (ratio of moles of solute to solvent) using the equation _

!

`a

`a

[O [O ln ⋯ [O 14


(16)

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where b and b are the chemical potential of pure solvent and solute respectively whereas A22 and B222 are the pair and triplet interaction term for solute particles. According to Hill,90 the ∗ coefficient A22 is related to the osmotic second virial coefficient through the relation ∗ c̅ X c 2

(17)

Here c and c̅ are the molecular volume of pure solvent and partial molecular volume of the ∗ solute at infinite dilution. X (= ) is the solute -solvent cluster integral.

For 1:1 electrolyte ⋯ 2 ln #∗

(18)

where #∗ is the nonelectrolyte contribution to the mean molal activity coefficient. From eqns (14) and (18), Y d /2 in which M1 is the molar mass of solvent in kg.mol−1. Using the relation, X c f.g O , eqn (16) can be written as ∗ h I

(19)

where g O is the isothermal compressibility of pure solvent and N is Avogadro constant. Thus using the values of Y obtained using eqn (14), the osmotic second virial coefficients were calculated for all the studied amino acid ionic liquids in aqueous solutions at 298.15 K and the data are reported in Table 7 along-with the corresponding Y and A22 values. ∗ The osmotic second virial coefficients (h ) are positive for all the studied ILs and their

magnitude increases with increase in hydrophobicity of ILs through the increase in alkyl chain ∗ length of amino acid anion on going from glycine to L-leucine. The positive values of h

indicates that repulsive type of ion-pair – ion-pair interactions dominates as the anion becomes 15



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∗ more hydrophobic. The linear relationship is observed between h and molar mass of the

alkyl chain-length (R) of anion as shown in Figure 8 which shows that hydrophobicity of ionic ∗ liquids is directly related with the osmotic second virial coefficient. In other words, the h can

be treated as a direct measure of hydrophobicity as revealed from the studied ionic liquids, however, large number of ILs and other similar systems needs to studied to make a generalized statement. d. Application of Pitzer Model According to Pitzer model,91-93 the osmotic coefficient for 1: 1 electrolyte can be expressed as 1 i M M j M

(20)

i M M k/ /T1 Xk/ V

(21)

where,

Here Aϕ is the Debye-Huckel constant for osmotic coefficient which is equal to 0.392 for aqueous solutions at 298.15 K, b is adjustable parameter having value of 1.2 kg1/2.mol−1/2 and k ∑ l is ionic strength of solution. The M and j M are the second virial and third virial type coefficient for the osmotic coefficient. j M coefficient is generally very small and is important in high concentrations of salt91 which we found negligibly small or even zero in some cases and hence not included in the present calculations. The coefficient M depends on concentration and is generally expressed as

M m m expqWk/ r 16


(22)

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m and m are Pitzer ion interaction parameters whereas W is the adjustable parameter having value of 2.0 kg1/2.mol−1/2. With the help of nonlinear least square fit method, the experimental osmotic coefficient data have been fitted using the eqns (20) to (22) to obtain the Pitzer ion interaction parameters m and m . The values obtained for m and m are included in Table 7. The standard deviation in osmotic coefficient obtained in the least square fit is also given in the last column of Table 7 and is calculated using the equation

s t

∑TMuEv UMwxy V !U4

/

z

(23)

where n is the number of data points and m is number of model parameters. The negative values of m suggest that the short-range interactions between unlike charged ions are dominant and the effect gets slightly reduced as the anion becomes more hydrophobic. This is reflected in the m values which increases with increase in hydrophobicity. As the interactions between like as well as unlike charged ions contribute to m , it is difficult to relate trends observed for m with hydrophobicity or any other similar parameter. It has been observed for many simple inorganic salts that the close correlation exits between m and m but in the present case no such correlation is observed between m and m . However, we observe significant correlation ∗ between osmotic second virial coefficient h obtained using McMillan-Mayer theory as above

and Pitzer ion interaction parameter m (see Figure 9) although both approaches are different. The exception for [Emim][Ile] ionic liquid can be attributed to the structural difference of [Ile]− ∗ ion. Thus, m can also be related to hydrophobicity like h as indicated earlier. Overall, all

these indicate that the interaction between like charges i.e. hydrophobic interaction (in present case anion-anion) are more prominent in aqueous solutions of [Emim][Leu] and [Emim][Ile] 17



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whereas interactions between unlike charges (cation-anion) are more significant in aqueous solutions of [Emim][Gly]. e. Mixing and excess thermodynamic properties for aqueous solutions of [Emim][Gly] Examination of Table 6 and Figure 4 shows that mixing process is enthalpy dominated as observed from the large negative values of mixing enthalpies. The mixing Gibbs free energy is slightly positive and is due to the major contribution of ideal mixing rather than the slight negative contribution arising from the nonideal part which is clearly reflected in Figure 5. Further, it can be seen from Figure 5 that all the excess thermodynamic properties are negative and the enthalpy-entropy effects are dominant over the Gibbs free energy changes. A major attention of researchers towards enthalpy-entropy compensation (EEC) effect is due to classical work of Lumary and Rajender94 highlighting the importance of this effect in biomolecular systems. Since then, many positive as well as negative arguments became the matter of debate on validity and usefulness of the EEC phenomenon.65-72,95 Cooper et al.66 have mentioned that the EEC is natural consequence of heat capacity changes and can arise due to limited free energy windows and quantum confinement effects. It has been noted earlier that the EEC effects are real and not just a result of some free energy windows, as they are the reflections of microheterogeniety in solution phase connected with cooperative structural changes of solvent water.52,96,97 Recent work of Starikov and Nordén69-72 proved that the EEC phenomenon is not only valid but is also not in conflict with conventional thermodynamics and can be derived in an exact way using statistical thermodynamics. We analysed the present system in accordance with the Starikov and Nordén model.69 The concentration dependent enthalpy-entropy compensation is observed for this system as shown in Figure 6a which obeys the linear relation ∆@ + 18


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.{|4} ∆ + X in which Tcomp is the compensation temperature and b has energy dimensions. The parameter b is found to negligible small (b = 0.15 J.mol−1) and the Tcomp is found to 307.54 K which slightly larger than the experimental temperature and this small temperature shift and the value of the parameter b is within the experimental error and cannot offer justifiable explanation about hidden Carnot cycle using the Starikov-Nordén model. This is due to the fact that the enthalpy and Gibbs free energy data used for EEC is expressed in joules per mole of solution meaning that the addition of small amount of solute (in the present case ~0 to 0.5 mol.kg−1) to large amount of solvent (55.51 mol.kg−1) does not affect the overall concentration change in the mixture as compared to that of pure water and hence the observed change in the corresponding thermodynamic properties if expressed per mole of solution instead of expressing them in per mole of solute. Similar observations have been reported earlier for aqueous solutions. Thus, after converting the enthalpy and Gibbs free energy data per mole of solute (the excess or mixing enthalpy per mole of solute is nothing but the molar enthalpy of solution), the observed enthalpyentropy compensation (Figure 6b) yields meaningful values of compensation temperature and parameter b. The values so obtained for .{|4} and for b are 224.32 K and -9.01 kJ.mol−1 respectively for the present system of aqueous [Emim][Gly]. These observations allows straightforward explanation of observed EEC using Starikov-Nordén model69 according to which if the EEC is a statistically significant and physically nontrivial correlation, the “hidden/imaginary Carnot process” may be represented by a set of microscopic heat engines, depending on the relation between the experimental temperature, T, and the compensation temperature, .{|4} ; namely, .{|4} > T (“heat pump”) and .{|4} < T (“refrigerator”). Thus, the much lower .{|4} than the experimental temperature for [Emim][Gly] + water system indicates that dynamical processes behind concentration increase can be viewed as “refrigerator”. Further, 19



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nonzero value for the b parameter suggests that mixing process involves the existence of microphases in equilibrium with each other leading to a cyclic sequence of reversible phase transitions (a closed thermodynamic cycle). This effect was also observed for aqueous solutions of amino acids.52 The entropic parameter (b/.{|4} ), also called as pertinent entropy difference, is found to be −40.16 J.mol−1 for aqueous [Emim][Gly] solutions clearly demonstrates the existence of microphases. The large negative values of excess entropies suggest that the hydration effects are more dominant along-with water structure promoting effects and this is due to the low charge density of these ions and the strong H-bonding ability of glycine anion which can helps to promote the cooperative H-bond formations favouring the water structure making effect (kosmotropic effect). This is being supported by the large negative value of entropic parameter (b/.{|4} ) as above using EEC analysis. This is further supported by the excess partial molar entropies of water (see Figure 7) which decreases with increase in concentration of ionic liquid. More interesting effect is observed in concentration dependent of excess partial molar entropies of solute, i.e. [Emim][Gly], which is negative and becomes less negative with increase in concentration of ionic liquid. The increase of partial molar entropies of solute indicates the existence of hydrophobic interactions in aqueous solutions of [Emim][Gly]. Since the glycine anion is not hydrophobic, the observed effect is due to the [Emim] cation suggesting that the existence of cation-cation hydrophobic interactions due to ethyl moiety of imidazolium ion. Thus water enforced hydrophobic association of low charge density cations further helps to promote the kosmotropic effect through cooperative H-bonding involving anion hydration rather than the dipole oriented ionic hydration as observed in ions with high charge density. 5. CONCLUSIONS 20


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Thermodynamic analysis of aqueous solutions 1-ethyl-3-methylimidazolium based amino acid ionic liquids at 298.15 K shows that these ionic liquids gets strongly hydrated and the corresponding high values of hydration numbers are due to the cooperative H-bonding of amino acid anions leading to the stabilization of water structure (kosmotropic effect). This is also supported strongly by negative entropic parameter obtained in EEC effect and also by concentration dependence of partial molal entropies of water, in aqueous [Emim][Gly] solutions, which decreases with increase in concentration of ionic liquid. The excess partial molar volume data is negative meaning that the hydrophobic hydration is important in aqueous solutions AAILs. The hydrophobic interaction also exists in aqueous AAIL solutions along-with strong cooperative anionic hydration. The anion-anion hydrophobic interaction are more prominent in aqueous solutions of [Emim][Leu] and [Emim][Ile] whereas interactions between unlike charges (cation-anion) are more significant in aqueous solutions of [Emim][Gly]. A linear correlation is observed between osmotic second virial coefficient and hydrophobicity of ionic liquids as well as Pitzer ion-ion interaction parameter m . Overall, this study and its extension for other amino acid ionic liquids with different cations will explore more information about ionic as well as noncovalent interactions which can help to understand potential use of these AAILs and it will also act as model systems towards better understanding of more complicated biological processes.

ACKNOWLEDGEMENTS

We greatly acknowledge the financial support (research project NO. SR/FT/CS-21/2011 with vide sanction number SERB/F/0694/2012-13) from Science and Engineering Research Board (SERB), New Delhi. K. R. Madkar and S.P. Shinde acknowledge UGC, New Delhi for financial assistance through SAP and BSR fellowship respectively. S.S. Barge acknowledges Department 21



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of Science and Technology (DST), New Delhi for financial assistance through Award of Inspire Fellowship.

AUTHOR INFORMATION Corresponding Author Email: [email protected]

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(79) Muhammad, N.; Man, Z. B.; Bustam, M. A.; Mutalib, M. I. A.; Wilfred, C. D.; Rafiq, S. Synthesis and thermophysical properties of low viscosity amino acid-based ionic liquids. J. Chem. Eng. Data 2011, 56, 3157-3162. (80) Tong, J.; Song, B.; Wang, C. -X.; Li, L.; Guan, W.; Fang, D. -W.; Yang, J. -Z. Prediction of the physicochemical properties of valine ionic liquids [Cnmim][Val] (n = 2, 3, 4, 5, 6) by semiempirical methods. Ind. Eng. Chem. Res. 2011, 50, 2418-2423. (81) Friedman, M. E.; Scheraga, H. A. Volume changes in hydrocarbon water systems. Partial molal volumes of alcohol water solutions. J. Phys. Chem. 1965, 69, 3795-3800. (82) Franks, F.; Smith, H. T. Apparent molal volumes and expansibilities of electrolytes in dilute aqueous solution. Trans. Faraday Soc. 1967, 63, 2586-2598. (83) Hagler, A. T.; Scheraga, H. A.; Nemethy, G. Current status of the water structure problem: Application to proteins. Annals. N. Y. Acad. Sci. 1973, 204, 51-78. (84) Franks, F. Hydrogen Bonded Solvent Systems, eds. A.K. Covington, and P. Jones, Taylor and Francis, London, 1968. (85) Wen, W. -Y.; Saito, S.; Lee, C. Activity and osmotic coefficients of four symmetrical tetraalkylammonium fluorides in aqueous solutions at 25 °C. J. Phys. Chem., 1966, 70, 12441248. (86) Burakowski, A.; Glinski, J. Hydration numbers of nonelectrolytes from acoustic methods. Chem. Rev. 2012, 112, 2059-2081. (87) Herrington, T. M.; Mole, E. L. Apparent molar volumes, temperatures of maximum density and osmotic coefficients of dilute aqueous hexamethylenetetramine solutions. J. Chem. Soc., Faraday Trans. 1, 1982, 78, 213-223.

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(88) Herrington, T. M.; Mole, E. L. Heats of dilution of aqueous solutions of sodium tetraphenylboron at 25 °C. J. Chem. Soc., Faraday Trans. 1, 1982, 78, 2095-2100. (89) McMillan, W.; Mayer, J. The statistical thermodynamics of multicomponent systems. J. Chem. Phys. 1945, 13, 276-305. (90) Hill, T. L. Theory of solutions. I. J. Am. Chem. Soc. 1957, 79, 4885-4890. (91) Pitzer, K. S.; Mayorga, G. Thermodynamics of electrolytes. II. Activity and osmotic coefficients for strong electrolytes with one or both ions univalent. J. Phys. Chem., 1973, 77, 2300-2308. (92) Pytkowicz, R. M. Activity Coefficients in Electrolyte Solutions, CRC Press, Boca Ranton, Florida, 1979. (93) Pitzer, K. S. Thermodynamics of electrolytes. I. Theoretical basis and general equations. J. Phys. Chem., 1973, 77, 268-277. (94) Lumry, R.; Rajender, S. Enthalpy–entropy compensation phenomena in water solutions of proteins and small molecules: A ubiquitous properly of water. Biopolymers 1970, 9, 11251227. (95) Cornish-Bowden, A. Enthalpy–entropy compensation: A phantom phenomenon. J. Biosci. 2002, 27, 121-126. (96) Leiter, H.; Patil, K. J.; Hertz, H. G. Search for hydrophobic association between small aprotic solutes from an application of the nuclear magnetic relaxation method. J. Solution Chem. 1983, 12, 503-517 (97) Holz, M.; Patil, K. J. Cation-cation association of tetramethylammonium ions in aqueous mixed electrolyte solutions. Ber. Bunsenges Phys. Chem. 1991, 95, 107-113.

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48


IIII) and Partial Molar Volume of Table 1. Density (Ρ), Apparent Molar Volume (φv ), Partial Molar Volume of Ionic Liquid ( IIII) for Aqueous Solutions of Amino Acid Ionic Liquids at 298.15 K Water ( m

ρ

10−3 ×

10−3 ×

10−3 ×

m

ρ

(mol.kg-1) (kg.m-3) (mm3.mol-1) (mm3.mol-1) (mm3.mol-1) (mol.kg-1) (kg.m-3) [Emim][Gly] +[H2O]

10−3 ×

10−3 ×

10−3 ×

(mm3.mol-1) (mm3.mol-1) (mm3.mol-1) [Emim][Leu] +[H2O]

0.00000

997.047

18.068

0.00000

997.047

18.068

0.01143

997.442 150.88 ±0.75 151.18

18.068

0.00881

997.328 209.92 ±0.95 210.07

18.068

0.02346

997.860 150.71 ±0.38 151.13

18.068

0.01878

997.644 210.01 ±0.42 210.21

18.068

0.04483

998.579 151.09 ±0.21 151.62

18.068

0.04142

998.337 210.50 ±0.19 210.74

18.068

0.08506

999.918 151.31 ±0.10 151.95

18.067

0.08475

999.657 210.59 ±0.10 210.83

18.068

0.10019

1000.393 151.59 ±0.09 152.25

18.067

0.10450

1000.251 210.60 ±0.08 210.83

18.068

0.14118

1001.693 151.89 ±0.08 152.56

18.066

0.15033

1001.613 210.59 ±0.06 210.78

18.067

0.21909

1004.117 152.16 ±0.05 152.77

18.066

0.16876

1002.131 210.73 ±0.06 210.90

18.068

0.27339

1005.771 152.26 ±0.05 152.78

18.065

0.19248

1002.820 210.73 ±0.05 210.85

18.068

0.33561

1007.659 152.27 ±0.05 152.63

18.066

0.21151

1003.383 210.64 ±0.05 210.74

18.068

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Page 34 of 57

0.39828

1009.445 152.49 ±0.04 152.66

18.067

0.26606

1004.926 210.66 ±0.05 210.66

18.068

0.44427

1010.830 152.38 ±0.05 152.40

18.068

0.32613

1006.580 210.70 ±0.04 210.57

18.069

0.51678

1012.920 152.38 ±0.04 152.11

18.070

0.38567

1008.220 210.61 ±0.04 210.35

18.070

0.57191

1014.517 152.30 ±0.04 151.80

18.073

0.45282

1009.964 210.69 ±0.04 210.27

18.071

0.49276

1011.067 210.53 ±0.04 210.01

18.073

0.56843

1013.009 210.50 ±0.02 209.78

18.075

[Emim][Ala] +[H2O] 0.00000

997.047

18.068

0.01617

997.578 166.75 ±0.81 166.91

18.068

0.02662

997.914 166.99 ±0.49 167.20

18.068

0.0000

997.047

18.068

0.04527

998.517 167.00 ±030 167.26

18.068

0.0090

997.351 207.92 ±1.19 208.03

18.068

0.09137

999.967 167.27 ±0.14 167.63

18.067

0.0183

997.669 207.72 ±0.59 207.91

18.068

0.12318

1000.951 167.37 ±0.11 167.76

18.067

0.0547

998.875 208.12 ±0.20 208.53

18.068

0.16387

1002.197 167.42 ±0.09 167.85

18.067

0.0720

999.438 208.22 ±0.15 208.69

18.067

0.18346

1002.780 167.50 ±0.08 167.93

18.067

0.0909

1000.027 208.51 ±0.13 209.03

18.067

0.23883

1004.428 167.57 ±0.07 167.99

18.066

0.1289

1001.210 208.75 ±0.09 209.31

18.067

0.29812

1006.124 167.74 ±0.06 168.12

18.066

0.1667

1002.388 208.76 ±0.08 209.29

18.066

0.36847

1008.129 167.78 ±0.05 168.08

18.066

0.2045

1003.520 208.92 ±0.07 209.35

18.066

[Emim][Ile] +[H2O]

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48


0.41870

1009.483 167.93 ±0.05 168.14

18.066

0.2429

1004.664 208.97 ±0.07 209.26

18.067

0.50125

1011.791 167.83 ±0.05 167.88

18.068

0.2846

1005.920 208.90 ±0.06 208.97

18.068

0.56143

1013.405 167.84 ±0.05 167.75

18.069

0.3359

1007.407 208.91 ±0.06 208.64

18.070

0.62514

1015.097 167.82 ±0.03 167.56

18.071

0.3821

1008.780 208.76 ±0.06 208.12

18.072

0.4303

1010.150 208.73 ±0.05 207.64

18.076

0.4856

1011.684 208.73 ±0.04 207.04

18.083

[Emim][Val] +[H2O] 0.00000

997.047

18.068

0.01193

997.371 200.60 ±0.54 200.77

18.068

0.02106

997.607 201.12 ±0.07 201.34

18.068

0.08831

999.344 201.36 ±0.06 201.64

18.068

0.10823

999.850 201.37 ±0.05 201.64

18.067

0.14975

1000.874 201.50 ±0.04 201.75

18.067

0.17707

1001.562 201.42 ±0.04 201.64

18.067

0.19581

1002.009 201.49 ±0.04 201.69

18.067

0.21432

1002.453 201.52 ±0.03 201.69

18.067

0.26824

1003.719 201.61 ±0.03 201.71

18.068

0.32944

1005.168 201.55 ±0.03 201.53

18.068 35



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48

0.39875

1006.742 201.56 ±0.03 201.42

18.069

0.44894

1007.876 201.53 ±0.03 201.28

18.070

0.51286

1009.286 201.50 ±0.03 201.12

18.072

0.57900

1010.736 201.43 ±0.03 200.89

18.074

36


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Page 37 of 57

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2. Limiting Apparent Molar Volume and Deviation Parameter for Amino Acid Ionic Liquids in Aqueous Solutions at 298.15 K

10−3 ×

10−9 × BV

(mm3.mol−1)

(mm6.mol−2)

[Emim][Gly]

150.4

5.52

[Emim][Ala]

166.5

1.44

[Emim][Val]

200.6

1.03

[Emim][Leu]

209.8

1.00

[Emim][Ile]

207.5

1.66

AAIL

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Page 38 of 57

Table 3. Osmotic Coefficient (φ), Water Activity (aw), Activity Coefficient of Water (γ1), Mean Molal Activity Coefficient (γ±) and Excess Gibbs Free Energy Change (∆GE) Data for Aqueous Solutions of Amino Acid Ionic Liquids at 298.15 K m

aW

γ1

γ±

(mol.kg−1)

∆GE

m

aW

γ1

γ±

(J.mol−1) (mol.kg−1) [Emim][Gly]+H2O

∆GE (J.mol−1)

[Emim][Leu]+H2O

0.01042 0.97188 0.99964 1.00001 0.89226

-0.16

0.00881 0.96210 0.99969 1.00001 0.90144

-0.10

0.02045 0.95743 0.99929 1.00003 0.85522

-0.41

0.01878 0.95293 0.99936 1.00003 0.86331

-0.34

0.04212 0.92975 0.99859 1.00011 0.80453

-1.10

0.04142 0.93736 0.99860 1.00009 0.81217

-1.08

0.08125 0.90586 0.99735 1.00027 0.74817

-2.81

0.08475 0.91626 0.99721 1.00025 0.75696

-2.96

0.10019 0.89973 0.99676 1.00036 0.72830

-3.83

0.10450 0.91429 0.99656 1.00032 0.73973

-4.05

0.14118 0.88956 0.99549 1.00055 0.69423

-6.32

0.15033 0.90954 0.99509 1.00048 0.70957

-6.82

0.17921 0.88540 0.99430 1.00072 0.66976

-9.00

0.16876 0.90026 0.99454 1.00059 0.70009

-7.79

0.21909 0.87837 0.99309 1.00093 0.64901 -11.91

0.19248 0.90506 0.99374 1.00063 0.68948

-9.57

0.27339 0.86938 0.99147 1.00124 0.62641 -16.06

0.21151 0.90025 0.99316 1.00073 0.68205 -10.75

0.33561 0.86010 0.98965 1.00162 0.60631 -20.99

0.26606 0.89584 0.99145 1.00095 0.66482 -14.54

38


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0.39828 0.85615 0.98779 1.00196 0.59067 -26.34

0.32613 0.89235 0.98957 1.00120 0.65101 -18.85

0.44427 0.86250 0.98629 1.00208 0.58151 -30.98

0.38567 0.89289 0.98767 1.00139 0.64108 -23.40

0.51678 0.86235 0.98407 1.00239 0.57034 -37.57

0.45282 0.89917 0.98544 1.00151 0.63320 -28.98

0.57191 0.86482 0.98234 1.00258 0.56413 -42.75

0.49276 0.90713 0.98402 1.00149 0.62983 -32.71

[Emim][Ala]+H2O

0.56843 0.91466 0.98144 1.00154 0.62560 -39.18

0.01048 0.96598 0.99964 1.00001 0.89216

-0.15

0.02024 0.95082 0.99931 1.00004 0.85641

-0.38

0.00958 0.96872 0.99967 1.00001 0.89815

-0.13

0.04183 0.93618 0.99859 1.00010 0.80659

-1.12

0.02180 0.94466 0.99926 1.00004 0.85552

-0.39

0.08286 0.91270 0.99728 1.00026 0.75005

-2.93

0.04684 0.92254 0.99844 1.00013 0.80512

-1.17

0.10445 0.90507 0.99660 1.00035 0.72919

-4.07

0.08782 0.91108 0.99712 1.00028 0.75728

-2.98

0.14735 0.89584 0.99526 1.00054 0.69722

-6.65

0.10001 0.90808 0.99673 1.00032 0.74696

-3.59

0.16893 0.89535 0.99457 1.00062 0.68439

-8.14

0.14775 0.90485 0.99519 1.00049 0.71591

-6.35

0.19420 0.89356 0.99377 1.00072 0.67133

-9.94

0.17052 0.90483 0.99446 1.00057 0.70474

-7.80

0.22182 0.88277 0.99297 1.00091 0.65901 -11.63

0.19090 0.90547 0.99379 1.00063 0.69613

-9.18

0.27333 0.87943 0.99138 1.00114 0.64012 -15.50

0.20329 0.90678 0.99338 1.00066 0.69145 -10.07

0.33515 0.87739 0.98946 1.00141 0.62263 -20.42

0.26811 0.90664 0.99128 1.00086 0.67203 -14.65

[Emim][Ile]+H2O

39



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48

0.39921 0.87699 0.98747 1.00167 0.60875 -25.74

0.31323 0.91078 0.98977 1.00094 0.66232 -18.17

0.45742 0.87684 0.98565 1.00190 0.59892 -30.65

0.41539 0.91029 0.98647 1.00123 0.64795 -25.71

0.52549 0.88017 0.98347 1.00209 0.59002 -36.70

0.43327 0.91011 0.98589 1.00128 0.64625 -27.01

0.57161 0.88122 0.98202 1.00224 0.58529 -40.68

0.50877 0.91243 0.98341 1.00144 0.64116 -32.66

[Emim][Val]+H2O

0.56567 0.91613 0.98150 1.00151 0.63915 -37.01

0.01017 0.96264 0.99965 1.00001 0.89409

-0.14

0.02109 0.94438 0.99928 1.00004 0.85482

-0.38

0.04419 0.93966 0.99850 1.00009 0.80428

-1.23

0.08831 0.91753 0.99708 1.00026 0.74799

-3.24

0.10823 0.90770 0.99647 1.00035 0.73027

-4.24

0.14975 0.90178 0.99515 1.00052 0.70155

-6.75

0.17707 0.89996 0.99428 1.00062 0.68679

-8.56

0.19581 0.89314 0.99372 1.00073 0.67804

-9.66

0.21432 0.88213 0.99321 1.00088 0.67029 -10.57 0.26824 0.88733 0.99146 1.00104 0.65177 -14.74 0.32944 0.88131 0.98959 1.00134 0.63614 -19.06 40


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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48


0.39875 0.88731 0.98733 1.00152 0.62331 -24.78 0.44894 0.89264 0.98567 1.00161 0.61643 -29.07 0.51286 0.89659 0.98357 1.00174 0.60997 -34.34 0.57900 0.89795 0.98144 1.00192 0.60547 -39.44

41



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Page 42 of 57

Table 4. Values of Coefficients Ai in Eqn (3) [Emim][Gly] Emim][Ala] Emim][Val] Emim][Leu] Emim][Ile] A2

0.315831

0.310697

0.336237

0.387818

0.415811

A3

-0.200100

-0.000321

0.007140

-0.030526

-0.045216

A4

0.206609

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 5. Hydration Number (h) and Distance of Closest Approach of Oppositely Charged Ions (a) for Amino Acid Ionic Liquids (AAIL) at 298.15 K AAIL

h

a (nm)

[Emim][Gly]

16.3

0.01

[Emim][Ala]

13.9

0.05

[Emim][Val]

15.1

0.05

[Emim][Leu]

13.5

0.09

[Emim][Ile]

13.5

0.11

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Page 44 of 57

Table 6. Thermodynamic Data for Aqueous Solutions of [Emim][Gly] at 298.15 K.a m

(∆Hmix or ∆HE)b,# ∆Hmix or ∆HE

A

T∆SE

∆Gmix

T∆Smix

( ̅ )E

( ̅ )E

(mol.kg−1)

(kJ.mol−1)

(J.mol−1)

(J.mol−1) (J.mol−1) (J.mol−1) (J.mol−1) (J.mol−1.K−1) (J.mol−1.K−1)

0.01057

-33.90

-12.9

440

-12.61

-13.8

0.89

-0.0002

-111.8

0.02126

-33.88

-25.9

460

-25.30

-27.5

1.56

-0.0005

-110.7

0.03074

-33.80

-37.4

540

-36.44

-39.4

2.02

-0.0011

-109.3

0.03898

-33.71

-47.3

630

-46.02

-49.6

2.33

-0.0020

-108.0

0.05018

-33.57

-60.6

770

-58.89

-63.2

2.65

-0.0037

-105.9

0.06025

-33.36

-72.3

980

-70.15

-75.2

2.99

-0.0060

-103.5

a

Note: The enthalpy, free-energy, and entropy values are reported for per mole of solution. Subscripts 1 and 2 correspond to solvent and solute, respectively. bValues are reported for per mole of solute. #Values from Ref. 76 reported as molar enthalpy of solution Hm, which is equal to molar enthalpy of mixing in joule per mole of solute.

44


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Table 7. Nonelectrolyte Solute-Solute Interaction Parameter (), Pair Interaction Terms for Solute Particles ( ), Osmotic Second Virial Coefficient (∗ ) and Pitzer Interaction Parameters for Amino Acid Ionic Liquids in Aqueous Solutions at 298.15 K AAIL

Y

A22

(kg.mol−1)

∗ 10−3 × h

m

m

s

(mm3.mol−1) (kg.mol−1) (kg.mol−1)

[Emim][Gly]

-0.18

-1.964

132.1

0.1116

-0.3682

0.0046

[Emim][Ala]

-0.10

-1.110

155.9

0.1225

-0.2710

0.0022

[Emim][Val]

-0.04

-0.467

195.8

0.1625

-0.3167

0.0044

[Emim][Leu]

0.02

0.190

211.0

0.1723

-0.2550

0.0035

[Emim][Ile]

0.06

0.653

212.8

0.1966

-0.2757

0.0057

45



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Page 46 of 57

Captions to Figures: Figure 1. The variation of ( √ ) as function of concentration, c, of ionic liquids in aqueous solutions amino acid ionic liquids at 298.15 K. [Emim][Gly], □; [Emim][Ala], ∆; [Emim][Val], ◊; [Emim][Leu], +; [Emim][Ile], ×. Figure 2. Osmotic coefficient, φ, for aqueous solutions of amino acid ionic liquids at 298.15 K. Debye-Huckel limiting law, ---; [Emim][Gly], □;

[Emim][Ala], ∆;

[Emim][Val], ◊;

[Emim][Leu], +; [Emim][Ile], ×. Figure 3. Mean molal activity coefficients (γ±) of amino acid ionic liquids in aqueous solutions at 298.15 K. [Emim][Gly],

; [Emim][Ala],

; [Emim][Val],

; Emim][Leu],

; [Emim][Ile], Figure 4. Mixing thermodynamic properties for aqueous solutions of [Emim][Gly] at 298.15 K. ∆Gmix, ; ∆Hmix, ; T∆Smix, . Figure 5. Excess thermodynamic properties for aqueous solutions of [Emim][Gly] at 298.15 K. ∆GE, ; ∆HE, ; T∆SE, . Figure 6. Correlation of excess enthalpy (∆HE) and excess entropy (∆SE) for aqueous solutions of [Emim][Gly] at 298.15 K. (a) ∆HE and ∆SE values are expressed in joule per mole of solution (b) ∆HE and ∆SE values are expressed in joule per mole of solute. Figure 7. Variation of excess partial molar entropies of water, < ̅ >+ , and of [Emim][Gly], < ̅ >+ , as a function of mole fraction of [Emim][Gly] at 298.15 K

46


Page 47 of 57

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 8. Correlation of osmotic second virial coefficient with molar mass of side chain (R) of amino acid anion of AAIL at 298.15 K. ∗ Figure 9. Correlation of osmotic second virial coefficient (h ) with Pitzer ion interaction

parameter (β(0)) for amino acid ionic liquids in aqueous solutions at 298.15 K

47


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

10−3 × (φV −SV × c1/2) /mm3.mol−1


Page 48 of 57

220 210 200 190 180 170 160 150 140 0

0.1 0.2 0.3 0.4 0.5 0.6 c /mol.L−1

Figure 1. The variation of ( √ ) as function of concentration, c, of ionic liquids in aqueous solutions amino acid ionic liquids at 298.15 K. [Emim][Gly], □; [Emim][Ala], ∆; [Emim][Val], ◊; [Emim][Leu], +; [Emim][Ile], ×.

48


Page 49 of 57

1.00 0.98 0.96 0.94 0.92

ϕ

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.90 0.88 0.86 0.84 0.0

0.2 0.4 0.6 . -1 ( m /mol kg )1/2

0.8

Figure 2. Osmotic coefficient, φ, for aqueous solutions of amino acid ionic liquids at 298.15 K. Debye-Huckel limiting law, ---; [Emim][Gly], □;

[Emim][Ala], ∆;

[Emim][Leu], +; [Emim][Ile], ×.

49


[Emim][Val], ◊;


0.0 -0.1 -0.2 lnγ±

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 50 of 57

-0.3 -0.4 -0.5 -0.6 -0.7 0.0

0.2

0.4

0.6

0.8

(m /mol.kg-1)1/2

Figure 3. Mean molal activity coefficients (γ±) of amino acid ionic liquids in aqueous solutions at 298.15 K. [Emim][Gly],

; [Emim][Ala],

; [Emim][Val],

; [Emim][Ile],

50


; Emim][Leu],

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


∆Gmix or ∆Hmix or T∆ ∆S mix /J.mol-1

Page 51 of 57

10 -10 -30 -50 -70 -90 0.000

0.001

0.002

0.003

x2 Figure 4. Mixing thermodynamic properties for aqueous solutions of [Emim][Gly] at 298.15 K. ∆Gmix, ; ∆Hmix, ; T∆Smix, .

51



0 ∆GE or ∆HE or T∆ ∆SE /J.mol-1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

-10 -20 -30 -40 -50 -60 -70 -80 0.000

0.001

0.002

0.003

x2 Figure 5. Excess thermodynamic properties for aqueous solutions of [Emim][Gly] at 298.15 K. ∆GE, ; ∆HE, ; T∆SE, .

52


Page 52 of 57

-0.30

∆SE / J.mol−1 -0.20 -0.10

∆HE = 307.54∆SE + 0.1507

-112

0.00 0

∆HE = 224.32∆SE - 9008.6

-10 -20 -30 -40 -50

∆SE / J.mol−1 -111 -110 -109

-108 -33.3 -33.4 -33.5 -33.6 -33.7 -33.8

-60 -70

-33.9

-80

-34.0

(a)

∆HE /kJ.mol− 1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


∆HE /J.mol− 1

Page 53 of 57

(b)

Figure 6. Correlation of excess enthalpy (∆HE) and excess entropy (∆SE) for aqueous solutions of [Emim][Gly] at 298.15 K. (a) ∆HE and ∆SE values are expressed in joule per mole of solution (b) ∆HE and ∆SE values are expressed in joule per mole of solute.

53



-95 -0.005

0 E 1

-100

)

2

-115

-0.010 -0.015

/J.mol−1.K−1

-105 -110

(S

0.000 1

E − S 20 /J.mol-1.K-1

-90

(S − S )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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-120 -0.020 0.000 0.001 0.002 0.003 x2 Figure 7. Variation of excess partial molar entropies of water, < ̅ >+ , and of [Emim][Gly], < ̅ >+ , as a function of mole fraction of [Emim][Gly] at 298.15 K

54


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


∗ h /(cm3.mol−1)

Page 55 of 57

220 210 200 190 180 170 160 150 140 130 120

[Emim][Leu] [Emim][Ile] [Emim][Val]

[Emim][Ala] [Emim][Gly] 0

20 40 60 . −1 Molar mass of R /(g mol )

Figure 8. Correlation of osmotic second virial coefficient with molar mass of side chain (R) of amino acid anion of AAIL at 298.15 K.

55



240

/(cm3.mol−1)

220

∗ h

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

[Emim][Leu] [Emim][Ile]

200

[Emim][Val]

180 160

[Emim][Ala]

140 [Emim][Gly] 120 100 0

0.1

0.2

0.3

β(0)/(kg.mol−1) ∗ Figure 9. Correlation of osmotic second virial coefficient (h ) with Pitzer ion interaction

parameter (β(0)) for amino acid ionic liquids in aqueous solutions at 298.15 K

56


Page 56 of 57

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Table of Content (TOC) Graphic

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Thermodynamic Studies of Ionic Hydration and Interactions for Amino

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