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Stability constants and thermodynamic functions of some rare earth metal ion chelates ofdl -α-aminobutyric acid
2076
Notes
<. 0
o
<~ 0
0
100
300
500
Reference temperature °C
Fig. 1. This inhibition has been confirmed by differential thermal analysis (DTA) of samples of AU containing differing amounts of nitrate. Samples were diluted with equal parts by volume of calcined alumina and heated at 5°C per minute in flowing dry argon. The resulting DTA traces, using alumina as a reference, are reproduced in Fig. 1. The peaks after 300°C are of interest. The exotherm at 400°C results from the self-reduction reaction and the endotherm between 500 and 60ff'C is caused by loss of oxygen as UO3 spontaneously decomposes to U3Os. The exotherms before 4000C are clearly ,associated with the NO3- present in the AU. Evidently, retained
ammonia is oxidised during these exothermic reactions, so that with increased nitrate content, pro~'essively less ammonia remains to effect self-reduction.
AAEC Research Establishment Lucas Heights, N.S. W. Australia
G. H. PRICE
REFERENCE 1. G. H. Price, J. Inorg. Nucl. Chem. 33, 4085 (1971).
J. inorg, nud, Chem. VoL 40, pp. 2076-2078 © Pergamon Press Ltd., 1978. Printed in Great Britain
0022-1902/7811201-20761502.00/0
S t a b i l i t y c o n s t a n t s a n d t h e r m o d y n a m i c f u n c t i o n s of some r a r e e a r t h m e t a l ion c h e l a t e s of D D - a - a m i n o b u t y r i c acid
(First received 19 May 1976; in revised form 29 March 1978) In earlier publications[I], we reported stability constants and thermodynamic functions of some metal chelates formed with DL-a-aminobutyric acid. Their formation was studied potentiometrically by Irving and Rossotti's method[2]. This note describes a similar study with some rare earth metal ions (La 3+, Ce 3+, Pr 3+, Nd 3+, Sm 3+ and y3+) and their stability constants and thermodynamic functions (AH, AG and AS) are reported.
Materials used All the rare earth metals were used as their nitrates, and their solutions were prepared in double distilled water, free from carbondioxide, containing a known concentration of nitric acid,
and were standardized by EDTA method[3]. Other reagents, apparatus and procedure, described earlier[l]. RESULTS AND DISCUSSION In these systems precipitation occurs from pH ~ 7 to 8, hence calculations have been done only up to those points before precipitation occurs. The formation curves (Fig. !, at 250C and 0.1 M ionic strength) show that in all cases ti values attain a maximum value of - I , before precipitation points, indicating that only one chelate ML is formed before hydrolysis sets in. The stability constants were hence computed by the average
Notes
2077
IO
05
0.0 4
5
6
pL Fig. 1. Formation curves for DL-a-aminobutyrate chelates. 0, La; O. Ce: A, Pr: A, Nd: IL Y: ~, Sm.
value method as t~ log ,-;-;------7;.,= log K + log [L]. The protonation constants as well as metal-ligand stability constants for the rare earth metal ion chelates, at four different temperatures (20, 25, 30 and 35°C) are reported in Table 1. in which the values of the standard deviation " ¢ " are also given. The results show a reasonably good fit of calculated and experimental curves. Thermodynamic parameters for various rare earth metal chelate systems are reported in Table 2. The AH values have been determined by temperature coefficient method, and the free energy and entropy changes were calculated as:
5O
,;5
AG = - RT In K and 4%9 2
AG = AH - TAS.
~ La
In Fig. 2, an attempt has been made to correlate log K values with their ionic radii[4]. It is observed that except for Y(III), rare earth metal ions show more or less a linear behaviour when their log K values are plotted against l/R[5]. This may be explained in terms of increasing coulombic attraction for the ligand with decreasin~cationic radius[6]. On this basis, the stability of Y(III) (R = 0.88A) chelate is expected to be much higher than lighter rare earth metal ions, whereas in the present case it is ~- Nd(III). Such observations have also been made earlier by Moeller et al. [7, 8] and have been explained in terms of the size of hydrated rare earth metal ions. It is argued that depending upon the nature
t
o~ .
~ Ce
f -
Pr
~
~1
Io4 Sm
Nd
I/R
Fig. 2. Correlation of log k values of OL-a-aminobutyrate chelates of some rare earth metal ions with their ionic radii.
of the ligand, the degree of solvation, and hence the electrostatic attraction for the ligand may vary. and thus this factor has an important bearing[9] in determining the stability of metal chelates. The results in Table 2, show low positive values of AH but high AS values, thus indicating that the rare earth a-aminobutyrate chelates are entropy stabilized. It is because, in aqueous
Table I. Stability constants of some rare-earth metal ion chelates formed with OL-a-aminobUtyric acid at different, temperatures (/~ = 0.1 M KNO3) Stability Constants
Metal ion Log K~H Log K2H La(lII) Ce(III) Pr(IlI) Nd(III) Sm(IIl) Y(lll)
pH range for calculations
Log K
tyt
Log K
tr
Log K
tr
Log K
o"
--6.6-7.9 6.3-7.5 6.0-7.3 6.0-7.2 5.8--6.9 5.8-7.1
9.63+0.02 2.34_+ 0.01 4.20-+0.03 4.62 _+0.02 4.79 _+0.02 4.92_+ 0.01 5.21 _+0.02 4.96_+0.02
--0.0060 0.0109 0.0114 0.0105 0.0118 0.0100
9.50_+0.02 2.30-+ 0.02 4.27-+0.02 4.67 _+0.02 4.86 _+0.01 5.01 -+0.03 5.26_+0.02 5.04_+0.03
--0.0089 0.0066 0.0061 0.0056 0.0097 0.0085
9,34_+0.03 2.25 -+0.01 4.35_+0.02 4.71 _+0.03 4.92 _+0.01 5.09_+0.02 5.32 _+0.03 5.11_+0.02
--0.0069 0.0085 0.0091 0.0054 0.0103 0.0049
9.19_+0.02 2.22 +-0.01 4.43_+0.02 4.76 _+0.03 4.98 _+0.02 5.17~.02 5.38 _+0.02 5.18_+0.02
--0.0108 0.0078 0.0084 0.0098 0.0080 0.0074
'o" = [number ~:£(Ar~)2 rvationsJ
20°C
25°C
] m , where Ar~ = tiexptl.- ticajea.
30oc
35oc
2078
Notes
Table 2. Thermodynamic parameters for some rare-earth metal ion chelates formed with OL-a-aminobutyric acid at 25°C and = 0.1 M(KNO3)
Metal ion La(III) Ce(lll) Pr(llI) Nd(llI) Sm(IIl) Y(III)
-AG (kcal/mole) 5.83 - 0.03 6.37-+0.03 6.63---0.01 6.84-+ 0.04 7.18 -+0.03 6.88 ± 0.04
AH (kcal/mole) 5.9 -+0.6 4.1 -+0.5 5.4-+0.4 6.6 ---0.5 4.3 - 0.5 6.2 -+0.5
Chemical Laboratories University of Allahabad Allahabad-211 002 India
AS (e.u.) 39 -+2 35 ± 2 40-+2 45 -+2 38 -+2 44 -4-2
solutions, complexation by chelation is favoured by release of bonded water molecules [9]. Moreover the net positive charge on the chelate species would also be less than that on the rare earth metal ion, thus resulting in a favourable entropy change.
tPostal address: Dr. M. N. Srivastava, 266, Mumfordganj (near Distillery), Allahabad-21l 002, India.
J. P. N. SRIVASTAVA M. N. SRIVASTAVA'~
REFERENCES
1. J. P. N. Srivastava and M. N. Srivastava, Ind. J. Chem. 14A, 818 (1976); Current Science 46, 443 (1977); Revue de Chimie Mindrale 14, 263 (1977). 2. H. Irving and H. S. Rossotti, J. Chem. Soc. 3397 (1953): 2904 (1954). 3. F. J. Welcher, The Analytical Uses of Ethylenediamine Tetraacetic Acid. p. 181. Van Nostrand, New York (1957). 4. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd Edn, p. 1052. Interscience, New York (1966). 5. T. Moeller, D. F. Martin, L. C. Thompson, R. Ferrus, G. R. Feistel and W. J. Randall. Chem. Rev. 65, 19 (1965). 6. J. C. Bailar, Jr. and H. J. Emeleus. Comprehensive Inorganic Chemistry, Vol. 4. pp. 28-36. Pergamon Press, New York (1973). 7. T. Moeller and L. C. Thompson. J. Inorg. Nacl. Chem. 24,499 (1962). 8. T. Moeller and R. Ferrus, lnorg. Chem. 1, 49 (1962). 9. I. Grenthe. Acta Chem. Scand. 17, 2487 (1963): 18, 283 (1964).
0022-1902/78/1~1-,'2078/$02.00/0
J. inorg, nucl. Chem. VoL 40. pp. 2078-2080 Pergamon Press Ltd.. 1978. Printed in Great Britain
Hydration of tri-n-hexylammonium picrate, tri-n-dodecylammonium picrate and picric acid in benzene solution (Received 15 March 1978) In connection with our studies on the hydration of halide ions of hydrogen bonded ion-pairs in benzene solution[I,2], we have determined the degree of hydration of the highly polarizable picrate ion in tertiary n-alkylammonium salts and in picric acid in benzene solution at 25°C, with the aim of elucidating the effect of anion polarizability on hydration and of determining the effect of water on the thermodynamic activity of the principal solute.
water content can therefore be represented by: mw = mw°aw + ~ms
(1)
EXI~TAL Tri-n-hexylammonium picrate (THxAHPi) and tri-n-dodecylammonium picrate (TLAHPi) were prepared from commercially available tri-n-hexylamine (pract., Fluka, A.G.) and tri-n-dodecylamine (purum, Rh6ne-Poulenc) as in Ref. [3]. Picric acid (HPi) (BDH) was purified as in Ref. [3]. Benzene (Riedel de Ha6n) was purified as in Ref. [4]. Density measurements. The density, d (g/ml), was measured in the same way as in Ref. [5]. Because no significant density difference between the dry and corresponding wet systems was detected, the density values of dry benzene solutions were used to convert molarities to molalities. Isopiestic method and analysis of water. See Ref. [1]. Vapour phase osmometry. See Ref. [5]. IR spectrophotometric measurements. See Ref. [2].
where mw° is the solubility of water in pure benzene at aw = 1 (the value of mw°=O.O393-+
RESULTS AND DISCUSSION
In as = In a* - haw
In spite of the fact that the solutes investigated undergo self-association as well as solvation in benzene solution[6] (the average association numbers in dry benzene are: 1.5 (ThxAHPi), 1.3 (TLAHPi) and 1.1 (HPi) at 25°C and a concentration of 0.15 moi/kg), the primary isopiestic data for THxAHPi, TLAHPi and HPi benzene solutions at 25°C, given in Table I, indicate, as in the benzene systems of the tertiary n-alkylammonium halides [1, 2], a linear dependence on water concentration, m~, and solute concentration, ms, at each water activity, aw. The
where a* is the thermodynamic activity of the solute in the absence of water. The effect of hydration on the thermodynamic activity of the solutes investigated is considerably lower than in the case of the tertiary n-alkylammonium halides in benzene solution. The h values of the picrate ion in tetra-n-butylammonium picrate in various solvents, calculated from the data given in Ref. [7], are 0.283-+ 0.008 (chlorobenzene), 0.276-+0.013 (o-dichloro-
(2)
Notes
<. 0
o
<~ 0
0
100
300
500
Reference temperature °C
Fig. 1. This inhibition has been confirmed by differential thermal analysis (DTA) of samples of AU containing differing amounts of nitrate. Samples were diluted with equal parts by volume of calcined alumina and heated at 5°C per minute in flowing dry argon. The resulting DTA traces, using alumina as a reference, are reproduced in Fig. 1. The peaks after 300°C are of interest. The exotherm at 400°C results from the self-reduction reaction and the endotherm between 500 and 60ff'C is caused by loss of oxygen as UO3 spontaneously decomposes to U3Os. The exotherms before 4000C are clearly ,associated with the NO3- present in the AU. Evidently, retained
ammonia is oxidised during these exothermic reactions, so that with increased nitrate content, pro~'essively less ammonia remains to effect self-reduction.
AAEC Research Establishment Lucas Heights, N.S. W. Australia
G. H. PRICE
REFERENCE 1. G. H. Price, J. Inorg. Nucl. Chem. 33, 4085 (1971).
J. inorg, nud, Chem. VoL 40, pp. 2076-2078 © Pergamon Press Ltd., 1978. Printed in Great Britain
0022-1902/7811201-20761502.00/0
S t a b i l i t y c o n s t a n t s a n d t h e r m o d y n a m i c f u n c t i o n s of some r a r e e a r t h m e t a l ion c h e l a t e s of D D - a - a m i n o b u t y r i c acid
(First received 19 May 1976; in revised form 29 March 1978) In earlier publications[I], we reported stability constants and thermodynamic functions of some metal chelates formed with DL-a-aminobutyric acid. Their formation was studied potentiometrically by Irving and Rossotti's method[2]. This note describes a similar study with some rare earth metal ions (La 3+, Ce 3+, Pr 3+, Nd 3+, Sm 3+ and y3+) and their stability constants and thermodynamic functions (AH, AG and AS) are reported.
Materials used All the rare earth metals were used as their nitrates, and their solutions were prepared in double distilled water, free from carbondioxide, containing a known concentration of nitric acid,
and were standardized by EDTA method[3]. Other reagents, apparatus and procedure, described earlier[l]. RESULTS AND DISCUSSION In these systems precipitation occurs from pH ~ 7 to 8, hence calculations have been done only up to those points before precipitation occurs. The formation curves (Fig. !, at 250C and 0.1 M ionic strength) show that in all cases ti values attain a maximum value of - I , before precipitation points, indicating that only one chelate ML is formed before hydrolysis sets in. The stability constants were hence computed by the average
Notes
2077
IO
05
0.0 4
5
6
pL Fig. 1. Formation curves for DL-a-aminobutyrate chelates. 0, La; O. Ce: A, Pr: A, Nd: IL Y: ~, Sm.
value method as t~ log ,-;-;------7;.,= log K + log [L]. The protonation constants as well as metal-ligand stability constants for the rare earth metal ion chelates, at four different temperatures (20, 25, 30 and 35°C) are reported in Table 1. in which the values of the standard deviation " ¢ " are also given. The results show a reasonably good fit of calculated and experimental curves. Thermodynamic parameters for various rare earth metal chelate systems are reported in Table 2. The AH values have been determined by temperature coefficient method, and the free energy and entropy changes were calculated as:
5O
,;5
AG = - RT In K and 4%9 2
AG = AH - TAS.
~ La
In Fig. 2, an attempt has been made to correlate log K values with their ionic radii[4]. It is observed that except for Y(III), rare earth metal ions show more or less a linear behaviour when their log K values are plotted against l/R[5]. This may be explained in terms of increasing coulombic attraction for the ligand with decreasin~cationic radius[6]. On this basis, the stability of Y(III) (R = 0.88A) chelate is expected to be much higher than lighter rare earth metal ions, whereas in the present case it is ~- Nd(III). Such observations have also been made earlier by Moeller et al. [7, 8] and have been explained in terms of the size of hydrated rare earth metal ions. It is argued that depending upon the nature
t
o~ .
~ Ce
f -
Pr
~
~1
Io4 Sm
Nd
I/R
Fig. 2. Correlation of log k values of OL-a-aminobutyrate chelates of some rare earth metal ions with their ionic radii.
of the ligand, the degree of solvation, and hence the electrostatic attraction for the ligand may vary. and thus this factor has an important bearing[9] in determining the stability of metal chelates. The results in Table 2, show low positive values of AH but high AS values, thus indicating that the rare earth a-aminobutyrate chelates are entropy stabilized. It is because, in aqueous
Table I. Stability constants of some rare-earth metal ion chelates formed with OL-a-aminobUtyric acid at different, temperatures (/~ = 0.1 M KNO3) Stability Constants
Metal ion Log K~H Log K2H La(lII) Ce(III) Pr(IlI) Nd(III) Sm(IIl) Y(lll)
pH range for calculations
Log K
tyt
Log K
tr
Log K
tr
Log K
o"
--6.6-7.9 6.3-7.5 6.0-7.3 6.0-7.2 5.8--6.9 5.8-7.1
9.63+0.02 2.34_+ 0.01 4.20-+0.03 4.62 _+0.02 4.79 _+0.02 4.92_+ 0.01 5.21 _+0.02 4.96_+0.02
--0.0060 0.0109 0.0114 0.0105 0.0118 0.0100
9.50_+0.02 2.30-+ 0.02 4.27-+0.02 4.67 _+0.02 4.86 _+0.01 5.01 -+0.03 5.26_+0.02 5.04_+0.03
--0.0089 0.0066 0.0061 0.0056 0.0097 0.0085
9,34_+0.03 2.25 -+0.01 4.35_+0.02 4.71 _+0.03 4.92 _+0.01 5.09_+0.02 5.32 _+0.03 5.11_+0.02
--0.0069 0.0085 0.0091 0.0054 0.0103 0.0049
9.19_+0.02 2.22 +-0.01 4.43_+0.02 4.76 _+0.03 4.98 _+0.02 5.17~.02 5.38 _+0.02 5.18_+0.02
--0.0108 0.0078 0.0084 0.0098 0.0080 0.0074
'o" = [number ~:£(Ar~)2 rvationsJ
20°C
25°C
] m , where Ar~ = tiexptl.- ticajea.
30oc
35oc
2078
Notes
Table 2. Thermodynamic parameters for some rare-earth metal ion chelates formed with OL-a-aminobutyric acid at 25°C and = 0.1 M(KNO3)
Metal ion La(III) Ce(lll) Pr(llI) Nd(llI) Sm(IIl) Y(III)
-AG (kcal/mole) 5.83 - 0.03 6.37-+0.03 6.63---0.01 6.84-+ 0.04 7.18 -+0.03 6.88 ± 0.04
AH (kcal/mole) 5.9 -+0.6 4.1 -+0.5 5.4-+0.4 6.6 ---0.5 4.3 - 0.5 6.2 -+0.5
Chemical Laboratories University of Allahabad Allahabad-211 002 India
AS (e.u.) 39 -+2 35 ± 2 40-+2 45 -+2 38 -+2 44 -4-2
solutions, complexation by chelation is favoured by release of bonded water molecules [9]. Moreover the net positive charge on the chelate species would also be less than that on the rare earth metal ion, thus resulting in a favourable entropy change.
tPostal address: Dr. M. N. Srivastava, 266, Mumfordganj (near Distillery), Allahabad-21l 002, India.
J. P. N. SRIVASTAVA M. N. SRIVASTAVA'~
REFERENCES
1. J. P. N. Srivastava and M. N. Srivastava, Ind. J. Chem. 14A, 818 (1976); Current Science 46, 443 (1977); Revue de Chimie Mindrale 14, 263 (1977). 2. H. Irving and H. S. Rossotti, J. Chem. Soc. 3397 (1953): 2904 (1954). 3. F. J. Welcher, The Analytical Uses of Ethylenediamine Tetraacetic Acid. p. 181. Van Nostrand, New York (1957). 4. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd Edn, p. 1052. Interscience, New York (1966). 5. T. Moeller, D. F. Martin, L. C. Thompson, R. Ferrus, G. R. Feistel and W. J. Randall. Chem. Rev. 65, 19 (1965). 6. J. C. Bailar, Jr. and H. J. Emeleus. Comprehensive Inorganic Chemistry, Vol. 4. pp. 28-36. Pergamon Press, New York (1973). 7. T. Moeller and L. C. Thompson. J. Inorg. Nacl. Chem. 24,499 (1962). 8. T. Moeller and R. Ferrus, lnorg. Chem. 1, 49 (1962). 9. I. Grenthe. Acta Chem. Scand. 17, 2487 (1963): 18, 283 (1964).
0022-1902/78/1~1-,'2078/$02.00/0
J. inorg, nucl. Chem. VoL 40. pp. 2078-2080 Pergamon Press Ltd.. 1978. Printed in Great Britain
Hydration of tri-n-hexylammonium picrate, tri-n-dodecylammonium picrate and picric acid in benzene solution (Received 15 March 1978) In connection with our studies on the hydration of halide ions of hydrogen bonded ion-pairs in benzene solution[I,2], we have determined the degree of hydration of the highly polarizable picrate ion in tertiary n-alkylammonium salts and in picric acid in benzene solution at 25°C, with the aim of elucidating the effect of anion polarizability on hydration and of determining the effect of water on the thermodynamic activity of the principal solute.
water content can therefore be represented by: mw = mw°aw + ~ms
(1)
EXI~TAL Tri-n-hexylammonium picrate (THxAHPi) and tri-n-dodecylammonium picrate (TLAHPi) were prepared from commercially available tri-n-hexylamine (pract., Fluka, A.G.) and tri-n-dodecylamine (purum, Rh6ne-Poulenc) as in Ref. [3]. Picric acid (HPi) (BDH) was purified as in Ref. [3]. Benzene (Riedel de Ha6n) was purified as in Ref. [4]. Density measurements. The density, d (g/ml), was measured in the same way as in Ref. [5]. Because no significant density difference between the dry and corresponding wet systems was detected, the density values of dry benzene solutions were used to convert molarities to molalities. Isopiestic method and analysis of water. See Ref. [1]. Vapour phase osmometry. See Ref. [5]. IR spectrophotometric measurements. See Ref. [2].
where mw° is the solubility of water in pure benzene at aw = 1 (the value of mw°=O.O393-+
RESULTS AND DISCUSSION
In as = In a* - haw
In spite of the fact that the solutes investigated undergo self-association as well as solvation in benzene solution[6] (the average association numbers in dry benzene are: 1.5 (ThxAHPi), 1.3 (TLAHPi) and 1.1 (HPi) at 25°C and a concentration of 0.15 moi/kg), the primary isopiestic data for THxAHPi, TLAHPi and HPi benzene solutions at 25°C, given in Table I, indicate, as in the benzene systems of the tertiary n-alkylammonium halides [1, 2], a linear dependence on water concentration, m~, and solute concentration, ms, at each water activity, aw. The
where a* is the thermodynamic activity of the solute in the absence of water. The effect of hydration on the thermodynamic activity of the solutes investigated is considerably lower than in the case of the tertiary n-alkylammonium halides in benzene solution. The h values of the picrate ion in tetra-n-butylammonium picrate in various solvents, calculated from the data given in Ref. [7], are 0.283-+ 0.008 (chlorobenzene), 0.276-+0.013 (o-dichloro-
(2)