Thermodynamics of ion-exchange on ferric antimonate

J inorg nucl. ('hem Vol. 43, pp

0022-1902/81/09212N08502.0010

2121-212~4. 1981

Pdn(ed in (;real Britain

Pergamon Press Ltd.

THERMODYNAMICS OF ION-EXCHANGE ON FERRIC ANTIMONATE J. P. RAWAT* and K. P. SINGH MUKTAWAT Department of Chemistry, Aligarh Muslim University, Aligarh-202001,India (Received 29 April 1980; received for publication 14 November 1980)

Abstract--A simple approach to ion-exchange equilibria on ferric antimonate has been applied. The values of selectivity coefficients for Ba2+, Mg2+, Ca2+ and Sr2÷ have been measured using equilibrium experiments at constant ionic strength and at different temperatures from 20 to 60°C. The thermodynamic equilibrium constant and values of AG°, AH° and AS° are reported. INTRODUCTION Inorganic ion-exchangers swell to a negligible extent and, therefore, the thermodynamics of ion-exchange on these materials will be simpler than on their organic counter parts where swelling is quite appreciable. In recent years several theories for ion-exchange equilibria have been developed and tried on inorganic ion-exchangers. Larsen studied the equilibria of Li+-H +, Na+-H + and K+-H + on zirconium phosphate[l;. Baetsle[2] studied tracer ion equilibria on zirconyl phosphate in hydrogen form using Rb', Cs ~, Sr 2", Ca 2' and Eu 3+. Nancollas[3] measured the equilibrium exchange of Li +, K + and Cs + ions with the hydrogen forms of semicrystalline zirconium phosphate at constant ionic strength. Ion-exchange isotherms for Li+-K + exchange were studied by Alberti[4] on crystalline zirconium phosphate. The present report describes the influence of temperature on the equilibria of MgZ+-Na +, BaZ+-Na +, Sr 2+Na + and Ca2+-Na + on ferric antimonate at constant ionic strength within the temperature range from 20-60°C. A simple approach has been applied and the thermodynamic parameters are calculated.

samples decreases as the hydrated radii of the ions decrease. This indicates that the metal ions enter the matrix in the hydrated form and the marked difference in the ion-exchange capacity is due to the difference in the hydration numbers of the ions. Similar behaviour has also been reported for zirconium phosphate [6]. Figures 1-4 show the reversibility of exchange of metal ions with the sodium form of ferric antimonate. The ion-exchange reaction may be written: 2 Na + + M2+~M 2+ + 2 Na +

EXPERIMENTAL

Ferric nitrate (B.D.H. India) and antimony pentachloride (B.D.H. England) were used. Antimony pentachloride was diluted with 4 M HCI to obtain the desired concentration. Ferric antimonate was synthesized as reported earlier[5;. Procedure. Water content of various cationic forms of the exchanger were determined by heating to a constant weight in an oven at 200°C for 2 hr. The exchange capacity was determined by eluting with 2 M sodium nitrate and 1.0 M alkaline earth chlorides or nitrates. The equilibrium experiments were performed by taking 20 ml solution containing sodium chloride or nitrate and the appropriate alkaline earth metal chloride or nitrate having a constant ionic strength 0.10. To this 0.2 of exchanger in sodium form was added and flasks were shaken in a temperature controlled S1CO shaker for 6 hr at the required temperature. Reversibility experiments were performed by equilibrating the exchanger in a particular metal ion form with that metal ion and sodium chloride solution at the different equilibrating temperatures. RESULTS AND DISCUSSION The ion-exchange capacity and the water content of the Mg, Ca, Sr and Ba forms of ferricantimonate are presented in Table 1. The water content in all four *Author to whom correspondence should be addressed.

The barred quantities referred to the exchanger phase and the unbarred to the solution phase. An examination of the exchange isotherms (Figs. I-4) shows that for exchange carried out at constant ionic strength, all ion-exchange isotherms are above the diagonal indicating that alkaline earth metals are preferred to sodium by ferric antimonate. The affinity for alkaline earth metals decreases as the temeprature increases. The isotherms indicate that uptake of barium ions occurs to a greater extent than of other alkaline earth ions. These results also show a small hystersis loop for BaZ+-Na + exchange while for Ca2+-Na +, SrZ+-Na + and Mg2+-Na + perfect reversibility is observed, These isotherms indicate the preference of alkaline earth metal ions is in order of Ba > Mg > Sr > Ca. The selectivity coefficients were calculated from the relationship (7]: 2

Kc

2

X M " X N a " TNa

X ~a" X ~ " vM

where )(M is the equivalent fractio,k of metal in the exchanger phase, XN, the equilivent fraction of sodium in the exchanger phase• XNa and XM mole fractions of sodium and metal in the solution phase respectively, yN~ and yM activity coefficients of sodium and metal ions respectively in the solution. The activity coefficient of the cations in the solution phase is calculated using the Debye-Huckel equation: -

•

log 7t

A~

V(~)

=

where A and B are constants, ai the ion size parameter and # the ionic strength. The values of A and B at the appropriate temperatures is taken from table given by Manov et al.[8] and the ion size, a, from that given by Freiser and Fernando[9]• 2121

Barium chloride

Ba 2+

nitrate 10.06

11.98

Strontium

S r 2+

ui t r a t e

12.5h

content

Ca l c i . m

'dater

C~'.2+

!

16.25

used

Na~nesillm chloride

Salt

Mg 2+

M~ta] ion

~

I

1.55

1.22

1.18

1.02

Ion-exch~e capacity of catiollic f o r u l o~ e x e h ~ I n g e r (meq/g)

!

1.56

1.23

1°20

1.02

Ion-exchauge capacity form of exchanger (raeq/g)

Table 1. Water contents and ion-exchange capacity of alkaline earth metal ions on Fe(III) antimonate of Na +

>

>

Z

E

>

Thermodynamics of ion-exchange on ferric antimonate

2123

I.O

o~ c

O.B-

L~

0.6-

x 20"C F o r w a r d 30"C o ~o*c • •

50C 60"C

•

2"0C B a c k w a r d

02kt~ I

0

I

0.2

t

0.4

I

0.6

I,O

08

Mole f r a c t i o n of Bo2÷ in s o l u t t o n

{ X Ba )

Fig. 1. Ion-exchange isotherm of Ba2+-Na + exchange on ferric antimonate.

1,0

,¢:

0.8

/// J ' J"

(b ¢: 0.6

0.4

•

/

c:

x

20"C B a c k w a r d 20*C For w a r d

*

3~c

o

4~C

6

5"0C

•

6~C

0.2 >. ~J

0 I/ 0

I

I 0.2

I

i 0.4

I

I 0.6

I

Mole f r a c t i o ~ of M g 2 * i n s o l u t ; o n

I 0.8

I 1.0

(XMg)

Fig. 2. Ion-exchange isotherm of MgZ+-Na + exchange on ferric antimonate.

JINC Vol, 43, No. 9K

2124

J.P. RAWAT and K. P. SINGH MUKTAWAT

• x

20C Backward 20°C Forward 30°C o 4(~C

U3

0.8 ¢: u

A

sEc

•

6(~C

.E " t~

0.6

o~ ~,o.4

o o~ 0 . 2

0~ 0

O. ;~

0.4

0.6

0.8

1.0

Mole fraction of Sr2*in solution (XSr ) Fig. 3. Ion-exchange isotherm of Sr:÷-Na ÷ exchange on ferric antimonate.

/

I,O

(J o.8-

0.6 -

x 20"C Forward * 3°°c o ~o'c • 50"C , 6~c • 2~C B o c k w a r d / / ~ t

/~ /1"i i'll x~./j" ..'X,~

~~-

(J

l.u

0,2! 0

.1

0.2

I

I

I

0.4

0.6

0.8

1.0

Ho(e frottl'on ofCo2*in 501ution (Xca) Fig. 4. Ion-exchange isotherm of Ca 2+ -Ha + exchange on ferric antimonate.

2125

Thermodynamics of ion-exchange on ferric antimonate

The selectivity coefficients are presented in Figs. 5-8. These results indicate that the value of Ko the selectivity coefficient does not remain constant but varies with the concentration of alkaline earth metal ions, XM, in the solid phase and hence K~ can be evaluated. Ko, the thermodynamic equilibrium constant is calculated from the expression given by Gains and

Thomas[lO] InKa = ( Z a - Z B ) +

/o'

lnKcdJ(M.

The values of K. are calculated with the help of plots of log Kc vs J?M and presented in Tables 2-5. These results

3.0

~

~ 4o'c

/

2.0

m o

x""

I.S

I,O

0.5

0

I

I

0

I

I

0.2

I

I

0.4

I

0.6

0.8

i.O

XBa

Fig. 5. Log of selectivity coefficients vs equivalent ionic fraction of Ba2+ in Na form exchanger.

2,8

•

20"C

• A

30C #,O'C

o

sSc

t

2.4

y

I/1, 2.0

~ o

1.6

1.2

~x~ x

~

x~'~

o ~ o ~

0.8

0.4

0 •0

I 0.1

I 0.2

i 0.3

I 0.4

I 0.5

I 0.6

I 0:7

P 0.8

I 0.9

1,0

#Mg

Fig. 6. Log of selectivity coe~cients vs equivalent ionic fraction of Mg2+ in Na form exchanger.

2126

J.P. RAWAT and K. P. SINGH MUKTAWAT

• 20=C o 3dc

~.O

o

_

,~

x

40C

,

5o'c

•

60"C

1.5

I.O

0.5

0

I

O

I

I

I

0.2

I

0.4

I

I

I

O.6

I

O.8

.O

XSr

Fig. 7. Log of selectivity coefficients vs equivalent ionic fraction of Srz+ in Ha form exchanger.

•

20"C

o 30"C

4o'c •

6dc

2.0

I.$

2/t

Q

,o

,0.5

i o l

o o

i 0.2

i 0.3

i 0.4

i o.s

i o.6

, *- To.~ o.B

~-'~ - ' ~ o~ ,.o

Ca

Fig. 8. Log of selectivity coefficients vs equivalent ionic fraction of Ca 2+ in Ha form exchanger. Table 2. Thermodynamic parameters for sodium-barium exchange on ferric antimonate at constant ionic strength t

TherBodynaBtc parameter

Ka AG"

(£J/mo le)

~, H" ( ] l ~ / B o l e ) ~- S° ( J / m o l e / d e E r e e )

w

l

!

T

20 * C

~0 * C

~0 * C

5 0 *C

88.65

56.98

32.01

19.66

9.71

-4.75

-~.~3

-3.92

-3.~0

-2,7~

-%7.052 -138,5

60 * C

2127

Thermodynamics of ion-exchange on ferric antimonate Table 3. Thermodynamic parameters for sodium-magnesium exchange on ferric antimonate at constant ionic strength t

K a A

G*

(KJ/mole)

A

H°

(KJ/mole)

.... !

• 20 C

Thermodynamic parm.eter

30

t

oC

40

'i

oc

!

o

50 C

60

0

C

13.15

10.10

8 +02

6.55

2.20

-2.73

-2.53

-2.36

-2.20

-0.95

-11.22

~--- S ° (J/mole/degree)

-28.92

Table 4. Thermodynamic parameters for sodium-strontium exchange on ferric antimonate at constant ionic strength ,

Thermodynamic

o

20

!

o

C

30

,

•

C

40

!

,

o

C

60°0

50 e

parameter

Ka

3.88

2.00

G° (KJ/mo l e )

-1.44

-o .76

H

o

1.31

I.

-o . ~

lO

1.Ol

-o .11

-o .o2

-lO.9O

(KJ/mole)

~ S ° (J/mole/degree)

-33.15

Table 5. Thermodynamic parameters for sodium-calcium exchange on antimonate at constant ionic strength !

Thermodynamic

! 20

•

C

! 30

•

!

C

40

0C

! 50

o

C

60

o

C

parameter

4.08

Ka o

" o

(~l,-oze)

-~ .49

2.15

~.69

1.56

1.18

-0.8~

-0.60

-0.52

-0.20

o

"~ rl

(KJl.,ole)

-17.28

A S"

(J/mole/degree)

-52.90

show that alkaline earth metals are preferred to Na ÷ form of ferric antimonate at all temperatures from 20 to 60°C. However, at higher temperatures the degree of selectivity decreases. The selectivity order of alkaline earth metals remains the same ( - supra), i.e. Ba > Mg > Sr > Ca. The standard free energy of exchange, AG O is calculated from the thermodynamic equilibrium constant using the equation AG ° -

RT ZN . Z

The entropy loss is an indication of greater order produced in the forward reaction when the alkaline earth metal ions is transferred to ferric antimonate. It reflects the difference in the solvation entropy of Na ÷ and the alkaline earth metal ions. Acknowledgements--The authors are grateful to Prof. W. Rahman for providing research facilities. We are thankful to Dr. O. P. Bansal for helpful suggestions and to C.S.I.R. (India) for financial assistance to one of us (K.P.S.M.).

ln K ,

where R is the gas constant and ZN,' ZM are the valencies of competing ionic species, T is the absolute temperature. (Tables 2-5). The standard enthalpy changes, AH ° and entropy changes, AS °, are presented in Tables 2-5.

REFERENCES

1. E. M. Larsen and D. R. Vissars, J. Chem. Phys. 64, 1732 (1960). 2. L. H. Baetsle, J. lnorg. NucL Chem. 25, 271 (1963), 3. G. H. Nancollas and B. V. K. S. R. A. Tilak, J. lnorg. NucL Chem. 31. 3643 (1969).

2128

J.P. RAWAT and K. P. SINGH MUKTAWAT

4. G. Alberti and U. Constantino, J. lnorg. Nucl. Chem. 36, 653 (1974). 5. J. P. Rawat and D. K. Singh, Anal. Chim. Acta 87, 157 (1976). 6. L. H. Baetsle, D. Huys and D. Van Deyak, J. lnorg. Nucl. Chem. 28, 2385 (1966). 7. R. M. Barrer and R. P. Townsend, J. C. S. Faraday II 72, 661 (1976).

8. G. G. Manov, R. G. Bates, W. J. Hamer and S. F. Acree, J. Am. Chem. Soc. 65, 1765 (1943). 9. H. Freiser and Q. Fernando, Ionic Equilibria in Analytical Chemistry. Wiley, New York (1966). I0. G. L. Gaines and H. C. Thomas, J. Chem. Phys. 21, 714

(1953).