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Thermodynamics of formation of hexoseborate complexes
J. inorg, nucl.Chem., 1970,Vol. 32, pp. 3545 to 3548. PergamonPress. Printed in Great Britain
T H E R M O D Y N A M I C S OF F O R M A T I O N OF H E X O S E BORATE COMPLEXES J. M. C O N N E R D e p a r t m e n t of C h e m i s t r y , Regis College, D e n v e r , Colo. 80221
(Received 6 April 1970) A b s t r a c t - T h e equilibrium c o n s t a n t s for complex formation b e t w e e n borate ion and the naturallyoccurring h e x o s e s were determined at four temperatures, and the corresponding t h e r m o d y n a m i c functions were calculated. Irregularities are explained on the basis of the complex structures INTRODUCTION
EQUILIBRIUM studies of complex formation b e t w e e n borate ion and a n u m b e r of organic p o l y h y d r o x y c o m p o u n d s were reported previously [ 1], with n u m e r o u s references to the prior literature. In this earlier work, all but three of the compounds were diols. T h e present study presents the results with the naturallyoccurring hexoses, which are pentols. Because of the relative inflexibility of the hexose lactoi structure, complex formation appears to occur only when hydroxyls on 1,2- or 1,3-carbons can orient in such a way that a 5- or 6 - m e m b e r e d ring can be f o r m e d with a tetrahedral borate ion, by the mutual loss of two molecules of water. In the case of fructose, complexing appears to occur largely with the m o r e flexible acyclic f o r m [2]. ---C--O
O--H-
/C) 0,,r t -\ -c _ _ o
--C--O
B
/
(C
\
O--H
~or i --C--O
1 : 1 complex anion
O
C--
B
/
(C) 0 ~,F
\
O - - C - -/
2 : 1 complex anion
Clearly this borate ion can simultaneously complex with a second molecule of hexose in the same way, to f o r m the 2 : 1 complex anion. T h u s we are concerned with the two equilibria B- + P ~ BP-
K~ --- [BP-] [B-][P]
and B-+2P
~ BP._,- K.z =
[BPe-] [B-][P] :~
where B- is borate ion and P is pentol (hexose). As in the past and as would be I. J. M. C o n n e r and V. C. Bulgrin, J. inorg, nucl. Chem. 29, 1953 (1967). 2. P.J. Antikainen. Suomen Kemistilehti 31B, 255 (1958). 3545
3546
J.M. C O N N E R
expected in a solvent of high dielectric constant, no evidence for a 1:2 complex was found. EXPERIMENTAL Stock solutions of the hexoses were prepared by weight from J. T. Baker reagent grade anhydrous D-hexose. After mutarotation was complete the concentrations were checked by polarimetry. The specific rotations of D-fructose, -- 92-4°; D-glucose, + 52'6°; and D-mannose, + 14.2° were given on the bottle labels. That of D-galactose, + 83'3 °, was obtained from the literature l3]. The stock solution of borax (sodium tetraborate decahydrate) was prepared from reagent grade N~2B4OT' 10 H20, and standardized by pH titration with standard HCI. Its aqueous solution is indistinguishable from a half-neutralized boric acid solution, and so is a good buffer. (The degree of complex formation is dependent on the pH.) All pH measurements were made with a Beckman Expanded Scale pH meter, equipped with a silver-silver chloride reference electrode. A Beckman Type E-2 glass electrode was used at the three upper temperatures and a Beckman No. 41263 glass electrode was used at 0°C. Both had silversilver chloride internals. The meter was standardized separately at each temperature, using the "borax buffer F" of Bates [4]. The constant temperature baths were set up and the reactions carried out as in the previous work[l]. The corrected temperatures of the baths were 0.00, 13.25, 24.95, and 35.01 +_-0.02°C. Since it is known that the equilibrium constants do not vary appreciably with borax concentration [1], only one concentration (0.02718 M borax) was used. Hexose concentrations were varied from 0-01 to 0-6 M. The highest corrected pH encountered was 9-49 for borax alone and the lowest was 4.69 for the most concentrated borax-fructose solution, both at 0°C. RESULTS AND DISCUSSION
It can be seen from Table 1 that the equilibrium constants lie in the same general range, except for fructose. Since the enthalpy changes for fructose complex formation are approximately the same as those for the other hexoses, the difference appears to be largely an entropy effect as shown in Table 2. Models of tile complexes were built, and three factors were found which should especially favor complex formation in the case of fructose. First, there were more (three) pairs of hydroxyl groups favorably oriented for borate complex formation with Table I. Equilibrium constants for formation of hexose-borate complexes
D- Fructose D-G alactose D-Glucose D-Mannose
0.00 °
13.25 °
24.95 °
K~
K~
Kj
K2
Kz
K.z
35.01 ° Kl K.e
5000 229,000 313 828 255 1240 107 1030
3800 214,000 240 530 205 1080 105 700
3000 124,000 175 428 150 760 102 458
1600 84,400 156 358 110 617 99 438
3. I. Heilbron et al., (Editors), Dictionary o f Organic Compounds 4th Edn. Vol. 3. p. 1487. Oxford University Press, New York (1965). 4. R.G. Bates,J. Res. natn. Bur. Stand. 66A, 179 (1962).
Formation
Table
of hexose-borate
2. Thermodynamics borate
D- F r u c t o s e
D-Galactose
D-Glucose D-Mannose
of formation
complexes
a(;o
complexes
~ 547
of hexose-
at 25 °
±H~,
~S,_°,
cal
cal
cal/deg
-- 4740
-- 3080
+ 5.6
-- 6950
-- 5060
+ 6.3
- - 3(160
-- 3060
0
- 3590
-- 3280
+ I .t1
- 2970
-- 4920
-- 6.5
- - 39311
-- 3650
+ 1.1)
- - 2741)
-- 357
-- 363(t
--4200
+ 8-(1 -- 1-9
the lactol form without bond strain or steric hindrance, than with the other three sugars. Second, in the case of the 2 : l complex, the chelate formed by using the hydroxyls of C~ and C3 of both fructose molecules leads to a very compact complex ion with seven opportunities for intraionic hydrogen bonding. These compact ions would be expected to be less solvated, and this is reflected in the relatively large positive entropy change for their formation. Third, there is evidence[2] that in aqueous solutions fructose exists mainly in the much more flexible acyclic form, and in this there are three pairs of hydroxyls which can become correctly oriented for a 5-membered chelate ring and similarly three pairs which can lead to a 6-membered ring. This is an unusually large number of possibilities. A word should be said here about the discrepancy in the case of the values in both Tables 1 and 2 for glucose, between the present work and that reported earlier[l]. In the previous work, considerable effort was made to keep mutaroration to a minimum before making the pH measurements. In view of the fact that the results obtained then showed no clear trend with temperature, in the present work all sugar solutions were allowed to achieve complete mutarotation before being mixed with borax. The observed entropy changes with glucose can be explained by two factors observed with models. For the [:1 complex ion, the best fit for a complexing borate occurs when the sugar molecule uses the hydroxyls of C4 and C6. This results in complete loss of rotational freedom for the C6-containing --CH=,OH group, and is reflected in the negative entropy change for formation of this complex. However, when a 2 : l complex forms, steric hindrance prevents the second glucose molecule from also complexing through the C,~ hydroxyl, and thus the second -CH2OH group is left free. Table 2 also shows a fairly large positive entropy change for the formation of the l : l mannose :borate complex. Models demonstrated that one of the best fits for a complexing borate ion involves the hydroxyls of C2 and C3. The resulting l : I complex is somewhat like the fructose complexes in being more compact than any possible galactose or glucose complex, with four opportunities for additional intraionic hydrogen bonding. Again such an ion would be expected to be less solvated, explaining the positive entropy change for its formation.
3548
J.M. C O N N E R
The method for calculation of the equilibrium constants and thermodynamic functions has been given before for the polyo! o-mannitol[1]. As was the case in the previous study, the observed pH changes with hexose concentration were corrected for the medium effect of dissolved hexose by a series of solutions of borax and non-complexing 2-methoxyethanol, equating [2-methoxyethanol] to [hexose].
T H E R M O D Y N A M I C S OF F O R M A T I O N OF H E X O S E BORATE COMPLEXES J. M. C O N N E R D e p a r t m e n t of C h e m i s t r y , Regis College, D e n v e r , Colo. 80221
(Received 6 April 1970) A b s t r a c t - T h e equilibrium c o n s t a n t s for complex formation b e t w e e n borate ion and the naturallyoccurring h e x o s e s were determined at four temperatures, and the corresponding t h e r m o d y n a m i c functions were calculated. Irregularities are explained on the basis of the complex structures INTRODUCTION
EQUILIBRIUM studies of complex formation b e t w e e n borate ion and a n u m b e r of organic p o l y h y d r o x y c o m p o u n d s were reported previously [ 1], with n u m e r o u s references to the prior literature. In this earlier work, all but three of the compounds were diols. T h e present study presents the results with the naturallyoccurring hexoses, which are pentols. Because of the relative inflexibility of the hexose lactoi structure, complex formation appears to occur only when hydroxyls on 1,2- or 1,3-carbons can orient in such a way that a 5- or 6 - m e m b e r e d ring can be f o r m e d with a tetrahedral borate ion, by the mutual loss of two molecules of water. In the case of fructose, complexing appears to occur largely with the m o r e flexible acyclic f o r m [2]. ---C--O
O--H-
/C) 0,,r t -\ -c _ _ o
--C--O
B
/
(C
\
O--H
~or i --C--O
1 : 1 complex anion
O
C--
B
/
(C) 0 ~,F
\
O - - C - -/
2 : 1 complex anion
Clearly this borate ion can simultaneously complex with a second molecule of hexose in the same way, to f o r m the 2 : 1 complex anion. T h u s we are concerned with the two equilibria B- + P ~ BP-
K~ --- [BP-] [B-][P]
and B-+2P
~ BP._,- K.z =
[BPe-] [B-][P] :~
where B- is borate ion and P is pentol (hexose). As in the past and as would be I. J. M. C o n n e r and V. C. Bulgrin, J. inorg, nucl. Chem. 29, 1953 (1967). 2. P.J. Antikainen. Suomen Kemistilehti 31B, 255 (1958). 3545
3546
J.M. C O N N E R
expected in a solvent of high dielectric constant, no evidence for a 1:2 complex was found. EXPERIMENTAL Stock solutions of the hexoses were prepared by weight from J. T. Baker reagent grade anhydrous D-hexose. After mutarotation was complete the concentrations were checked by polarimetry. The specific rotations of D-fructose, -- 92-4°; D-glucose, + 52'6°; and D-mannose, + 14.2° were given on the bottle labels. That of D-galactose, + 83'3 °, was obtained from the literature l3]. The stock solution of borax (sodium tetraborate decahydrate) was prepared from reagent grade N~2B4OT' 10 H20, and standardized by pH titration with standard HCI. Its aqueous solution is indistinguishable from a half-neutralized boric acid solution, and so is a good buffer. (The degree of complex formation is dependent on the pH.) All pH measurements were made with a Beckman Expanded Scale pH meter, equipped with a silver-silver chloride reference electrode. A Beckman Type E-2 glass electrode was used at the three upper temperatures and a Beckman No. 41263 glass electrode was used at 0°C. Both had silversilver chloride internals. The meter was standardized separately at each temperature, using the "borax buffer F" of Bates [4]. The constant temperature baths were set up and the reactions carried out as in the previous work[l]. The corrected temperatures of the baths were 0.00, 13.25, 24.95, and 35.01 +_-0.02°C. Since it is known that the equilibrium constants do not vary appreciably with borax concentration [1], only one concentration (0.02718 M borax) was used. Hexose concentrations were varied from 0-01 to 0-6 M. The highest corrected pH encountered was 9-49 for borax alone and the lowest was 4.69 for the most concentrated borax-fructose solution, both at 0°C. RESULTS AND DISCUSSION
It can be seen from Table 1 that the equilibrium constants lie in the same general range, except for fructose. Since the enthalpy changes for fructose complex formation are approximately the same as those for the other hexoses, the difference appears to be largely an entropy effect as shown in Table 2. Models of tile complexes were built, and three factors were found which should especially favor complex formation in the case of fructose. First, there were more (three) pairs of hydroxyl groups favorably oriented for borate complex formation with Table I. Equilibrium constants for formation of hexose-borate complexes
D- Fructose D-G alactose D-Glucose D-Mannose
0.00 °
13.25 °
24.95 °
K~
K~
Kj
K2
Kz
K.z
35.01 ° Kl K.e
5000 229,000 313 828 255 1240 107 1030
3800 214,000 240 530 205 1080 105 700
3000 124,000 175 428 150 760 102 458
1600 84,400 156 358 110 617 99 438
3. I. Heilbron et al., (Editors), Dictionary o f Organic Compounds 4th Edn. Vol. 3. p. 1487. Oxford University Press, New York (1965). 4. R.G. Bates,J. Res. natn. Bur. Stand. 66A, 179 (1962).
Formation
Table
of hexose-borate
2. Thermodynamics borate
D- F r u c t o s e
D-Galactose
D-Glucose D-Mannose
of formation
complexes
a(;o
complexes
~ 547
of hexose-
at 25 °
±H~,
~S,_°,
cal
cal
cal/deg
-- 4740
-- 3080
+ 5.6
-- 6950
-- 5060
+ 6.3
- - 3(160
-- 3060
0
- 3590
-- 3280
+ I .t1
- 2970
-- 4920
-- 6.5
- - 39311
-- 3650
+ 1.1)
- - 2741)
-- 357
-- 363(t
--4200
+ 8-(1 -- 1-9
the lactol form without bond strain or steric hindrance, than with the other three sugars. Second, in the case of the 2 : l complex, the chelate formed by using the hydroxyls of C~ and C3 of both fructose molecules leads to a very compact complex ion with seven opportunities for intraionic hydrogen bonding. These compact ions would be expected to be less solvated, and this is reflected in the relatively large positive entropy change for their formation. Third, there is evidence[2] that in aqueous solutions fructose exists mainly in the much more flexible acyclic form, and in this there are three pairs of hydroxyls which can become correctly oriented for a 5-membered chelate ring and similarly three pairs which can lead to a 6-membered ring. This is an unusually large number of possibilities. A word should be said here about the discrepancy in the case of the values in both Tables 1 and 2 for glucose, between the present work and that reported earlier[l]. In the previous work, considerable effort was made to keep mutaroration to a minimum before making the pH measurements. In view of the fact that the results obtained then showed no clear trend with temperature, in the present work all sugar solutions were allowed to achieve complete mutarotation before being mixed with borax. The observed entropy changes with glucose can be explained by two factors observed with models. For the [:1 complex ion, the best fit for a complexing borate occurs when the sugar molecule uses the hydroxyls of C4 and C6. This results in complete loss of rotational freedom for the C6-containing --CH=,OH group, and is reflected in the negative entropy change for formation of this complex. However, when a 2 : l complex forms, steric hindrance prevents the second glucose molecule from also complexing through the C,~ hydroxyl, and thus the second -CH2OH group is left free. Table 2 also shows a fairly large positive entropy change for the formation of the l : l mannose :borate complex. Models demonstrated that one of the best fits for a complexing borate ion involves the hydroxyls of C2 and C3. The resulting l : I complex is somewhat like the fructose complexes in being more compact than any possible galactose or glucose complex, with four opportunities for additional intraionic hydrogen bonding. Again such an ion would be expected to be less solvated, explaining the positive entropy change for its formation.
3548
J.M. C O N N E R
The method for calculation of the equilibrium constants and thermodynamic functions has been given before for the polyo! o-mannitol[1]. As was the case in the previous study, the observed pH changes with hexose concentration were corrected for the medium effect of dissolved hexose by a series of solutions of borax and non-complexing 2-methoxyethanol, equating [2-methoxyethanol] to [hexose].