Indian Journal of Chemistry Vol. 39A, June 2000, pp. 598- 602

Partial molar volumes and viscosities of some monovalent ions in ethanolamine and water--ethanolamine mixtures at 298.15 K

(Late) R K Wadi, Vinita & Rita Kakkar*t

Department of Chemistry, University of Delhi , Delhi II 0 001·

Received 22 June 1999; revised 22 October 1999

Precise densities and viscosi ties for lithium chloride, potassium chloride, potassium iodide, cesium iodide and sodium tetra phenyl borate in ethanolamine and water-ethanolamine mixtures (I 0, 20, 30, 40, 50, 60, 70, 80 and 90 mass% ethanolamine) upto a maximum salt molality of 0.2 mol kg·1 are reported at 25"C. The electrolyte apparent mol ar volumes have been calcu lated and extrapolated to infinite dilution using the Masson equation to obtain the limiting electrolyte partial

molar volumes. The limiting ionic partial mol ar volumes, V '\.m have been estimated using Mukerjee's method . Variations

in<!>;," with change in solvent composition are discussed in terms of the changes in the solvent structure. The experimental results for viscosity have been analyzed using the Jones-Dole equation. The 8-coefficient values have been obtained and their vari ations are discussed with changes in the solvent structure.

Introduction It is well-known that the reaction medium plays an

important role in determining reactivity, which is reflected in thermodynamic, transport and spectral properties 1.2. In order to gain insight into the mechanism of such interactions, thermodynamic and transport studies involving one or more solutes in mixed solvent systems are highly useful. The volumetric and viscometric studies of electrolytes at infinite dilution in various mixed solvent systems have contributed to our knowledge about electrolyte­non-electrolyte-water interactions. By examining the viscosity B-coefficient and <l>v0 of ions as a function of size, nature, temperature and composition of the mixed solvent, it is possible to study the effect of these parameters on ion-water interactions, with the hope of obtaining a better understanding of the interactions in solutions.

In continuation of our earlier work3'4 on the

volumetric and conductometric behavior of solutions of a few alka li halides in ethanolamine and the aqueous ethanolamine system, we report herein the same for lithium and potassi um chlorides, potassium and cesium iodides and sodium tetraphenyl borate.

Materials and Methods Ethanolamine (S .D. Fine-chem., A.R., 99%) was

E-mai l: rita_kakkar@vsn l.com

further purified over 4 A molecular sieves for more than' 48 h, and then by fractional distillation through a one-meter column, in a stream of nitrogen under reduced pressure. The middle fraction of the distillate was collected in an all glass system, and then sealed with parafilm and stored in a dessicator over P20 5.

Prior to use, freshly distilled ethanolamine samples were degassified in an all set vacuum line. The water content, tested by the Karl Fischer titration instrument (Orion Carl Fischer Titrator Model AF-8 by Orion Research Incprporation), was always found to be less than 0.1 % for the experimental samples, which were earlier found by gas chromatography (SE 30 column) to contain no significant amounts of impurities.

Doubly distilled water was deionized by passing through two Cole-Parmer mixed bed ion-exchanger col umns (conductivity <8X 10·7 S-cm·\ KCI (Sarabhai M Chemicals, G.R.) was precipitated from a saturated conductivity water solution by adding an equal amount of reagent grade ethyl alcohol. Crystals were dried at 363 K in a vacuum oven and kept in vacuo over P20 5 for 24 h before use . LiCI (Sisco, A.R.), Csi (Sigma, USA, AR) and KI (Merck, GR) were dried at 363 K and kept in vacuo over P20s before use. NaBPh4 (Sigma, USA, AR) was dried at 323 K and kept in vacuo over P20 5 before use. The solutions were prepared by mass in Borosil glass bottles and all transfers of the solutions were done

WADI et al.: PARTIAL MOLAR VOLUMES & VISCOSITIES OF MONOVALENT IONS 599

under nitrogen atmosphere in a glove box to prevent contact with atmospheric C02•

The densities of solutions were determined with an Anton Paar vibrating tube digital densitometer (DMA 601/60). The U-tube of the densitometer was dried and flushed with nitrogen before injecting the experi­mental solution. The densities of water (p=0.997054 g cm-3

) and dry nitrogen (p=1.1456x10·3

g em·\ used for the calibration of the instrument, were taken from the literature5

·6

. The temperature of the measuring cell was controlled to 0.01 °C by circulating water from a thermostat. Ttre temperature inside the densitometer cell was monitored with a calibrated Anton Paar DT-100-20 digi tal thermometer with an accuracy of ±0.01 K. The reproducibility of the density values was of the order of ±10 ppm. The partial molar volume of sodium chloride, found to be 16.66 cm3 mor 1 in water, agreed well with the reported values 7 .

For viscosity measurements, the concentration of the solutions varied in the range of 0.01-0.2 mol r 1

and the solutions were prepared by weight dilution of the stock sol ution . The conversion of molality into molarity was done using the density values . Measurements were made using a Ubbelohde suspended level viscometer. All measurements were done in a water thermostat having an accuracy of ± 0.01 K. The flow time was noted down to a precision of ± 0.05 s using a Fischer stop watch of least count 0.0 I s. The absolute viscosity of pure ethanolamine agreed well with the values reported in the literature8

.

Results and Discussion The apparent molar volumes (AMV), <l>v. were

calculated from solution densities (p=pb and p1) using Eq. (1)

<l>v=M s1 p- { IOOO(p -Po)/ (ms P Po) } ... (1)

where M., ms and Po are the solute molar mass, solute molality and the solvent density respectively. The concentration c on the molar scale were calculated using Eq. (2)

c=[(IOOO ms Pr)/ ( IOOO+M s ms)] . .. (2)

The densities, Po. for the water-ethanolamine binary mixtures are recorded in Table l along with the

d. 1· I 9 10 0 I . correspon mg tterature va ues · . ur va ues are m fairly good agreement with the reported values. The densities, Pr. of solutions of LiCl , KI, KCl, Csl and NaBPh4, in various water-ethanolamine mixtures and pure ethanolamine, for salt concentrations ranging from 0.01-0.2 mol dm·3 at 25°C, are reported in ref. 11. The electrolytes' apparent molar volumes were calculated using Eq . 1, where Po is the density of the binary solvent and p is the densi ty of the ternary mixture. The density is found to be a linear function of molality, i.e.

... (3)

The regression coefficients thus obtained are reported in ref. II.

The electrolytes ' AMV s, <l>v. in ternary solutions were calculated from Eq. 1, and are reported along

Table !-Densi ties of water-ethanolamine mixtures and the apparent mol ar vol ume of ethanolamine at 298.15 K.

J cm·3

Mass% EA XEA .. This work Ref. 9 Ref. 3 <l>v

0(EA)

h

10 0.0317 1.0007 1.0007 59.61

20 0.0687 1.0052 1.0053 59.09

30 0. 1122 1.0 I 03 1.0103 1.0 I 03 58.58

40 0.1642 1.0157 1.0157 58.09

50 0.2278 I 0206 1.0207 1.0207 57 .82

tiO 0.3067 1.0243 I 0261 1.0244 57.71

70 0.4077 1.0259 1.0255 1.0259 57.88

80 0.54 12 1.0247 1.0255 1.0259 58.34

85 0.6256 1.0220 1.0224 58.80

90 0.7264 10197 1.02 14 1.0197 59.20

100 1.0000 1.0117 1011 7 10 117 60.29c

aMole fraction EA hU nits: cm3mor 1 cref. I 0

600 INDIAN J CHEM, SEC. A, JUNE 2000

with the density data in Table I. Application of the Redlich-Meyer equation 12 was not possible due to the lack of data on the compressibility and pressure variation of dielectric constant, necessary to calculate the theoretical limiting slope, Sv *. Thus we used the Masson equation

. .. (4)

where Sv is the experimental slope. Within the experimental error, our values for <l>v0 varied linearly with --./c. Uncertainties in <l>v0 values were estimated to be of the order of ± 0.3 cm3 mor 1

• <l>v o for KI was compared with the literature3

, and was found to be in agreement.

The ionic partial molar volumes, V 0 ;0 n, were calculated using Mukerjee 's empirical method 13

• As a

first approximation, V\ was taken as zero; the values

of V0;011 based on thi s were then plotted against the

cube of the Pauling crystal radi i, r/ The V\ was then adjusted until the best fit for both the cations and anions yielded the same linear plot.

70 • Gsl

6 8 o I<I

66 0. KCI

64 /-·~ /. . 62 ./· .\ 60 • 58 • 56

T 54 -0 E 52

I')

E " 50 ....

!ft. 48 o_.....o-o-o~

_.,if

46 0/0 0

44

42

40

30

28

26

24L---L----L--~----L-~ 0 20 40 60 80 100

Mdss% OF EA

Fig. !-Limiting partial molar volume of electrolytes against

1 mass% of ethanolamine

V o tr for different ions from water to the water­ethanolamine mixture reflects the interactions between the dissolved electrolyte and the solvent.

Values for V o tr were calculated using the relation

V 0 tr= V 0

ms-V 0 w, ... (5)

where V0ms and V0

w are the partial molar volumes of the electrolyte in the mixed solvent and in water, respectively.

The <l>v 0 versus mass% ethanolamine curve (Fig. I) shows a maximum between 60-70 mass% ethanolamine for KCI, KI and Csl. A similar trend was observed for NaCl, NaBr and Nai in the same

system3. The increase in <l>v 0 and V 0 1, for KCl , KI and

Csi till 60-70 mass% ethanolamine may be due to a decrease in electrostriction in the presence of ethanolamine. Thus the electrostriction effect, which brings about shrinkage in the volume of the solvent, is decreased in mixed solvents as compared to that in pure water. Since electrostriction primarily reflects electrolyte-solvent interactions, it can be inferred that the electrolyte-solvent interaction increases with

280~-----------------------,

272

T 0 268 E

266

0 20 I 40 GO

e NoBPit4'

0 Li CI

80

Mass% OF EA

100

Fig. 2-Limiting partial molar volumes of electrolytes against

I mass% of ethanolamine

WADI eta/.: PARTIAL MOLAR VOLUMES & VISCOSITIES OF MONOVALENT IONS 601

increase in ethanolamine till 60-70 mass% ethanolamine, and is maximum in this region. Further

the decrease in the value of <l>v 0 and V 0 tr in going

from 70 mass% ethanolamine to pure ethanolamine, may be explained as due to strong solvent-solvent interactions3

'14.

For LiCl and NaBPh4 (Fig. 2); the decrease in <l>v0

or negative V01r in going from pure water to pure

ethanolamime may be due to the small size of lithium and the hydrophobic behavior of BPh4-, because of which the interactions among the solvent components become strong.

The data 11 for relative viscosities of the electrolytes in various water-ethanolamine mixtures, measured at 298.15 K, were analyzed with the Jones­Dole equation

(lllllo-1)/ ...fc=A-B ...fc ... (6)

where 11 and llo are the viscosities of the ternary solutions and binary solvent, respectively, and cis the molar concentration. Best estimates of the Jones-Dole B-parameter were found using the theoretical values for A and then calculating the B-parameter. The A coefficients were calculated at 298.15 K from the physical properties of the solvent using the Falkenhagen - Vernon equation,

A=0.2577 Aot (llo (Et)A0.5 Ao + A0-)

[1- 0.6863 (Ao+- 1--..o-IAoi] ... (7)

The conductance data required in these calculations have been taken from ref. 11 and the dielectric constant data from ref. 4. The A-coefficients, thus calculated, are reported in Table 2. The viscosity B­coefficients, so calculated using the least squares method, are presented in Table 3.

For LiCl and NaBPh4, the B-coefficients are found to be positive and increase with increase in ethanolamine content (Fig. 3), indicating the existence of strong ion-solvent interactions. The magnitude of the B-coefficient also increases, indicating the simultaneous increase in ion-solvent interactions. In other words, the solvation of ions increases as the ethanolamine content increases in water. Thus these electrolytes are structure makers for

the solvent mixture structure. The negative V 0 rr

values for these electrolytes also support the above behavior. For KCl, KI and Csl, initially the B­coefficient is negative and decreases till 40-50 mass% ethanolamine (Fig. 3); it then starts increasing and becomes positive after 60 mass% ethanolamine. These minima in B-coefficient values occur in the range of the mixed solvent composition corresponding to the maximum stabilization of the water structure by ethanolamine molecules . In other words, minima may be connected to the increase in the structure breaking effect of these electrolytes due to stabilization of the water structure by ethanolamine molecules . After 60 mass% ethanolamine, the positive values of the B-coefficient indicate strong alignment of the solvent molecules in the vicinity of the ions, which promotes the solvent mixture structure. Thus it suggests the structure making effect of these electrolytes in the ethanolamine rich region

and similar behavior is supported by V 0 1r values.

Based on the sign of ( V 0 ion-V 0 crys)7 for the studied

ions in various water-ethanolamine mixtures, Lt and Na+ appear to be structure makers, K+ a borderl ine case, and c s·, cr, r and BPh4- are structure breakers of the solvent-mixture structure.

Table 2-Theoretical A-coefficients of the Jones-Dole equ ation for electrolytes at 298.15 K

Mass% EA 103 A/dm312mol' 112

LiCl KCI KI Csl NaBPh4

0 6.81 5.2 1 5.01 4.89 12.23

10 6.96 4.68 5.20 4.93 11.96

20 6.75 4.58 4.82 4.93 11.42

30 6.49 4.54 4.72 4.68 12.16

40 6.51 4.49 5.27 4.65 10.51

50 6.58 4.23 4.67 3.75 9.37

60 7.0 1 3.91 4.48 4.42 8.46

70 7.59 4.03 4.51 4.67 6.33

80 8.25 4.83 5.45 4.97 10.82

90 10.39 5.53 6.25 5.54 1390

100 13.27 6.18 7.35 6.49 13.36

602 INDIAN J CHEM, SEC. A, JUNE 2000

Table 3- Viscosity B-coefficient of electrolytes in water+ethanolarnine mixtures at 298. 15 K

Mass% EA LiCl KCl

10 -0.141 -0.022

(0.010) (0.002)

20 0.153 -0.030

(0.001) (0.001)

30 0.160 -0.022

(0.003) (0.002)

40 0.199 -0.012

(0.001) (0.001)

50 0.186 0.020

(0.002) (0.001)

60 0.241 0.101

(0.005) (0.00 1)

70 0.314 0.19 1

(0.002) (0.002)

80 0.450 0.351

(0.003) (0.005)

85

90 0.770 0.671

(0.002) (0.003)

100 1.089 0.902

(0.002) (0.002)

~

I 160 0 E

...., 140 E "0

12 0 ..... IXl

100 )(

N

~ 80

60

40

20

0

-2 0 0 20 40 60 ao 100

Mass% OF EA

Fig. 3- Viscosity 8 -coefficients of electrolytes versus mass% ethanolamine. ( I) CsCI , (2) Kl, (3) KCl , (4) LiCl , (5) NaBPh4

B/dm3 mor 1

KI Csl NaBPh4

0.079 -0.141 1.239

(0.002) (0.007) (0.002)

-0.087 -0.161 1.225

(0.003) (0.001 ) (0.009)

-0.096 -0.175 1.309

(0.001) (0.002) (0.003)

-0.102 -0.153 1.424

(0.002) (0.001) (0.007)

-0.068 -0.135 1.489

(0.002) (0.003) (0.004)

-0.1017 -0.049 1.549

(0.008) (0.002) (0.002)

0.167 0.130 1.61 9

(0.002) (0.002) (0.003)

0.336 1.691

(0.003) (0.002)

0.362

(0.004)

0.61 9 1.731

(0.002) (0.002)

0.862 0.740 1.751

(0.003) (0.002) (0.002)

Acknowledgement The award of a research fellowship to Vinita from

UGC, New Delhi is gratefully acknowledged.

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