AQUEOUS INFRARED STUDIES OF
THE
PYRIDINE CARBOXYLIC ACIDS
2153
Aqueous Infrared Studies of the Pyridine Carboxylic Acids by J. F. Wojcikl and T. H.Stock Department of Chemistry, Villanova University, Villanova, Pennsylvania 10086
(Received July 1, 1068)
The infrared spectra of the pyridine carboxylic acids have been studied in DzO in the region 1760-1350 cm-1. The spectra are obtained as a function of added strong acid and base. Some spectra were also run in dioxane. Bands are assigned and on the basis of the position of these bands, their behavior upon acidification and neutraliEation, and their behavior in dioxane, a dipolar structure is assigned to all three acids in aqueous solution. A correlation of the dipolar species/neutral species ratio with water concentration in water-dioxane mixtures is also presented. The reasoning which led to the assignments of the Introduction bands is as follows: The band assigned to the carbonylEvidence for the dipolar (D in Figure 1) nature of the stretching motion (V+COOH) appears only in solution to pyridine carboxylic acids in water has been presented which strong acid has been added. Under this condias a result of studies of their ultraviolet ~ p e c t r a . ~ , ~ tion both the nitrogen and the carboxyl group will be Some aqueous infrared work has been done,4 but this protonated. This band appears in the region expected work has led to the speculation that the acids exist as for this type of motion. The ammonium carboxylate nondipolar species with internal hydrogen bonding. antisymmetric stretch (V+COO -) appears in the region Martell, et aLj6and Sawyer, et U Z . , ~ have made several A plot of the intensity of this previously noted.6N6 studies on aqueous solutions of amino acids which led band against the concentration of the monoprotonated to the following conclusions. An un-ionized carboxyl 3 for nicotinic acid. The conacid is given in Figure group absorbs around 1700-1730 cm-l, ammonium centrations of monoprotonated acid were calculated carboxylate absorbs around 1620-1630 cm-l, and an from the amounts of DC1 or NaOD added to the a-amino carboxylate absorbs around 1575-1585 cm-’. organic acid using the pK’s of nicotinic acid in H207 The shift of the carboxylate band to higher frequencies D20 by means of the data in Be1L8 and corrected to when adjacent to a positively charged nitrogen was The correlation is somewhat affected by the fact that attributed to an inductive effect. The results prethis band is adjacent to a band attributed to ring vibrasented here support that explanation. tion which does not diminish in basic solution. Thus, Experimental Section the contribution from the ring band makes the intensity The pyridine carboxylic acids were purchased from of the plotted band somewhat higher in basic solution. Eastman Organic Chemicals and were purified by The band attributed to an antisymmetric carboxylate sublimation. No attempt was made to deuterate these stretching motion without the adjacent positive charge compounds. DC1 (12 M) was purchased from Merck ( V C O O -) behaves as expected. This band disappears as Sharp and Dohme and D2O (99.7 mol %) was purthe basicity of the solution is lowered. The position of chased from General Dynamics. Dioxane was Mathethis band is somewhat lower than expected, but it is son Coleman and Bell Spectroquality reagent. A the only band in this frequency range that has the sodium deuterioxide solution was prepared by dissolvexpected intensity behavior. ing in D2O metallic sodium enclosed in glass tubing. The assignment for the bands below 1400 cm-l (HThe DC1 and NaOD solutions were standardized in the bonded ring deformation) is based on a comparison with usual manner. All solutions were prepared by weighing results obtained from the spectra of carboxylic acids in the organic acid into a 2-ml volumetric flask, adding the standardized acid or base from a microburet and dilut(1) Author to whom correspondence should be addressed. ing with solvent. (2) R. W. Green and H. K. Tong, J . Amer. Chem. SOC.,78, 4896 Infrared spectra were measured on a Perkin-Elmer (1956). Model 337 grating spectrophotometer. Barnes Engi(3) H.P. Stephenson and H. Sponer, ibid., 79, 2050 (1957). neering calcium fluoride cells with adjustable path (4) M. Paris, G. Thomas, and J. C. Merlin, Bull. SOC.Chim. Fr., 707 (1961). lengths were used. All spectra were calibrated with (5) K.Nakamoto, Y.Morimoto, and A. E. Martell, J . Amer. Chem. polystyrene film. Soc., 85, 309 (1963). Results and Discussion (6) D.T.Sawyer and J. E. Tackett, ibid., 85, 314 (1963). (7) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” The results for the DzO solution studies of the three Butterworth and Co. Ltd., London, 1959,p 535. acids are given in Tables I, 11, and 111. Representa(8) R.P. Bell, “The Proton in Chemistry,” Cornel1 University Press, tive spectra are given in Figure 2. Ithaca, N. Y.,1959,p 189. Volume 73, Number 7 July 1969
2154
J. F. WOJCIKAND T. H. STOCK
Table I : Infrared Spectra of Picolinic Acid" in DtO Moles of DC1
or NaOD per mole of HA
1-00 (DC1) 0.70 (DC1) 0.50 (DC1) 0.20 (DC1) 0.00 0.20 (NaOD) 0.50 (NaOD) 0.70 (NaOD) 1.00 (NaOD)
v+COOH
v+coo-
1730 (s) 1730 (s) 1730 (m) 1724 (w)
1650 (8) 1650 (s) 1650 (s) 1650 (s) 1650 (s) 1650 (s) 1650 ( 6 ) 1650 (w)
...
...
...
...
Arom. ring def.
Arom. ring def.
1613 (8) 1610 (s) 1613 (s) 1619 (s) 1610 ( 6 ) 1610 ( 6 ) 1608 ( 6 ) 1608 (s) 1608 (s)
1587 (m) 1587 (m) 1587 (m) 1587 (m) 1587 ( 8 ) 1587 (s) 1585 (s) 1585 (s) 1582 (s)
...
vcoo
... ... ...
...
1567 (w) 1565 (8) 1565 (s) 1565 (s)
Arom. ring def.
Arom. ring def.
...
1453 (m) 1451 (m) 1451 (m) 1451 (m) 1451 (m) 1449 (w)
1502 (w) 1502 (w) 1499 (w) 1504 (m) 1502 (w)
...
...
...
...
... ...
Hbonded ring def.
1376 (s) 1376 (s) 1376 (s) 1376 (s) 1377 (s) 1379 (s) 1385 (s) 1389 ( 6 ) 1389 (s)
a Picolinic acid concentration is 1.0 M . Ionic strength is 1.0 and maintained with KC1. All spectra run at room temperature. Cell thickness is 0.015 mm.
Table II: Infrared Spectra of Nicotinic Acid" Moles of DCI or NaOD per mole of HA
1.OO (DCl) 0.71 (DCl) 0.51 (DCl) 0.21 (DC1) 0.00 0.20 (NaOD) 0.49 (NaOD) 0.69 (NaOD) 0.98 (NaOD)
+COOH
vtcoo-
1715 (s) 1715 ( 6 ) 1715 (e.) 1715 (m)
1634 (w) 1634 (m) 1634 (s) 1634 (s) 1634 (s) 1634 (8) 1634 (s) 1634 (m)
... ... ... ... ...
...
Arom. ring def.
Arom.
1613 (w) 1605 (m) 1605 ( 8 ) 1605 (e) 1605 ( 8 ) 1605 (s) 1605 (s) 1600 (s) 1600 (s)
1575 (w) 1577 (w) 1577 (m) 1580 (8) 1580 (s) 1580 (s) 1587 (sh)
ring def.
... ...
woo-
...
... , ,
,
...
...
Arom. ring def.
Arom. ring def.
1495 (w) 1493 (w) 1493 (w) 1493 (w) 1493 (w)
1453 (m) 1451 (m) 1451 (m) 1449 (m) 1447 (m)
... ... ... ...
1564 (ah) 1563 (sh) 1560 (m) 1558 (m)
... ... ... ...
Hbonded ring def.
1379 (sh) 1379 (s) 1379 (s) 1379 (s) 1379 (s) 1381 (s) 1385 (s) 1385 (s) 1387 (a)
Hbonded ring def.
1359 (8) 1369 (s) 1359 (sh)
... ... ...
... ...
...
" Nicotinic acid concentrations were 0.162 M . Ionic strength is 0.162 and maintained with KCl. All spectra run at room temperature. Cell thickness is 0.1 mm.
Table 111: Infrared Spectra of Isonicotinic Acid in Dz0 Mol- of DC1 or NaOD per mole of H A
1.00 (DC1) 0.70 (DCl) 0.50 (DCl) 0.21 (DCl) 0.00 0.20 (NaOD) 0.50 (NaOD) 0.70 (NaOD) 1.00 (NaOD)
-
V+COOH
v+coo
1721 (s) 1721 (s) 1721 (m) 1709 (w)
1613 (s) 1613 (8) 1613 (s) 1616 (8) 1616 (s) 1616 (s) 1613 (sh)
.,.
...
... ... ...
... ...
Arom ring def.
... ... ..,
...
... 1608 (sh) 1600 (s) 1595 (s) 1590 (s)
Arom.
Arom. ring def.
woo-
...
...
1575 (w) 1575 (w) 1575 (w) 1575 (w) 1575 (w) 1575 (sh) 1582 (sh) 1582 (sh)
...
... ...
... 1453 (w) 1543 (w) 1543 (m) 1543 (m)
ring def.
1458 ( 6 ) 1458 (8) 1458 (s) 1458 (m) 1458 (w)
... ... ... ...
H-
H-
bonded ring def.
bonded ring def.
1397 (sh) 1397 (s) 1397 (s) 1397 (s) 1381 (s) 1383 (s) 1387 (s) 1387 (s)
1355 (9) 1361 (sh)
1389 (s)
... .
I
.
* . .
... ...
...
...
" Isonicotinic acid concentration is 0.077 M . Ionic strength is 0.077 and maintained with KC1. All spectra run at room temperature. Cell thickness is 0.1 mm.
nonpolar solvent^.^ For deuterated acids, bands appear in this frequency region which are attributed to ring vibrations involving acid dimers. In D2O one can the conceive Of a hydrogen-bonded acid and D2O molecules which forms a ring. The posiThe Journal of Physical Chemiatry
tion of this band should be sensitive to whether or not the carboxyl group is protonated. The observed (9) C. N. R.Rao, "Chemical Applications of Infrared Spectroscopy," Academic Press, Inc., New York, N. Y., 1963, p 222.
AQUEOUS INFRARED STUDIES OF
THE
2155
PYRIDINE CARBOXYLIC ACIDS
A
/
0.5 0.4
s C
0
5 0.2 8
z D
3
A-
Figure 1. Equilibria pertinent to the system.
[HA], mol/l. Figure 3. Dependence of absorbance of v+coo- on the concentration of HA ([HA] = [N] [ D ] ) : 0 , solutions with DCl added; H, no DC1 or NaOD added; A, solutions with NaOD added.
+
I
/
3.c
6,O
7.0
Wavelength, 1.1. Figure 2. Representative spectra: curve A, picolinic acid; curve B, isonicotinic acid. Both spectra run in DzO with no strong acid or base. See Tables I and I11 for details. [The instrument is linear in wavelength. All readings were converted into wave numbers in the text.]
bands are somewhat SO. I n addition the authors have observed bands in this region for simple aliphatic amino acids in addition to weak bands which can be attributed to -CH2- scissoring motion.1° The remaining peaks are attributed to ring vibrations (aromatic ring deformation) This assignment is based on the known spectra of similar compounds. For instance, methyl picolinate has bands at 1590, 1575, 1475, and 1433 cm-l.ll The intensities of the bands obtained here are sensitive to the degree of protonation of the acid.
1.C 0.425
0.450
0975
Q - $/@Df 11 Figure 4. Dependence of K d on solvent dielectric constant. Point taken from value of Kd obtained from extrapolation in Figure 5 .
Based on the location of the bands involving motion of the carboxyl group and their behavior upon acidification or neutralization of the solutions of the acids, a dipolar structure is assigned to all three acids. It is noted that the shift of the carboxylate band ( V + C O O - ) (10) J. F. Wojcik and T. Stock, to be published. (11) A. R. Katritsky and A. R. Hands, J . Chem. Soc., 2202 (1958).
Volume 79, Number 7 July 1969
J. F. WOJCIKAND T. H. STOCK
2156
to higher frequencies is achieved in the cases where the protonated nitrogen is not adjacent to the carboxyl group. I n these cases the ring is a good transmitter of the positive charge. The values obtained for the ratio of dipolar molecule to nondipolar form in aqueous solution suggested a further analysis of the system. The value of this ratio for picolinic, nicotinic, and isonicotinic acids is 15, 10, and 25, respectively.2 Simple electrostatics predict that if these values were valid, the nondipolar form would exist in a solution of the acid in a solvent of low dielectric constant. This would show up in the infrared spectra by a loss of v+coo- and the appearance of vCOOn. Such a measurement was made in dioxane (D
2.c
.
G
m
VCOOH
was absent in the spectrum of the acid in the DgO
so that it was concluded that only the dipolar form
existed in DzO. Using the intensity of V+COO- in DzO and the concentration of the acid, the molar absorptivity of the dipolar species was estimated. This mola,r absorptivity was then used to estimate the dipolar species concentration in the DzO-dioxane mixtures. The plot is not linear at low dielectric constants. The slope for the line drawn at high dielectric constant is 46. This corresponds to a p 2 / u 3 value of 1.9 X 10-l2. If a is assigned a reasonable value such as 4 A, then a value of 11 D is obtained for the dipole moment of the dipolar species in D20-dioxane mixtures. This compares favorably with the value of 12 D estimated for glycine.18 The nonlinearity of the plot is not surprising. A recent publication has suggested that the dielectric constants is not important but that one should include in the equilibrium constant the water concentration raised to a power reflecting the degree of solvation of the various species involved in the eq~i1ibrium.l~Such a constant is demonstrated to be independent of solvent composition for water-dioxane mixtures. The complete constant for the case studied here is given by
0 1.5 3. N
which can be rewritten as 1.c
A plot of log Kd US. log C H z O (in units of molarity) should give a straight line with slope equal to --n. This plot is given in Figure 5 and is in fact linear. The slope of 2.7 suggests that the dipolar species is more Figure 5. Dependence of Kd on water concentration.
= 2.21), and the prediction was borne out for picolinic and nicotinic acids. VOOOH for the picolinic acid is split into two bands at 1761 and 1730 cm-I. For nicotinic acid this same band appears at 1730 cm-l. The solubility of isonicotinic acid was too low in dioxane to obtain useful spectra. Based on equations derived by Kirkwood,l 2 one concludes that a plot of log Kd (Kd being defined in Figure 1) os. ( D - 1)/(2D l),where D is the dielectric constant of the solvent, should be a straight line. The slope of the line is predicted to be equal to p 2 / u a k T , where p is the dipole moment of the dipolar molecule and a is a molecular diameter. Such a plot is given in Figure 4 for picolinic acid. Kd was obtained in the following manner: V+COO- is characteristic of the dipolar species. It was assumed that the concentration of the dipolar molecule and the intensity of the peak were related by Beer’s law.
+
The Journal of Physical Chemistry
highly solvated than the neutral species. One can also extrapolate to a water concentration corresponding to pure water. The value of Kd determined in this manner for pure water is 14.6. This agrees very well with the value of 15 reported earlier.2 The results obtained here present one more class of equilibria with a complete equilibrium constant independent of solvent composition, and while it is not clear that an explanation neglecting the dielectric properties of the solvent is valid, the analysis seems to provide a good extrapolation method. It was noted that V C O O H for picolinic acid is in fact split into two bands. Over the concentration range 0.02-0.29 M , the ratio of the intensities of these two bands remains constant at 2.1 f 0.1. This suggests an equilibrium in dioxane of the sort (12) J. G.Kirkwood, J . Chem. Phys., 2, 361 (1934). (13) J. P. Greenstein and M. Winitz, “Chemistry of the Amino Acids,” Vol. 1, John Wiley & Sons, New York, N. Y.,1961,p 468. (14) A. 8. Quist and W. L. Marshall, J . Phys. Chem., 72, 1636 (1968).
SOLIDSOLUTION OF ALUMINUM OXIDEIN RUTILETITANIUM DIOXIDE
0 r
F=O H-0
- Qc=o
I
\O-H.*a
I1
It has previously been proposed that I represented picolinic acid in D20.4 Internal hydrogen bonding of this sort was also proposed for nicotinic and isonicotinic acids. It is difficult to see how an internal hydrogen
2157
bond can be formed in the latter two cases due to the separation of the hydrogen bonding atoms. As the data in this work show, the spectra can be explained in terms of the dipolar species existing in DzO. In summary, by assuming the dipolar form for the three acids as the principal species in aqueous solution, a variety of solution infrared results can be explained.
Acknowledgment. The authors wish to thank the American Chemical Society for supporting this work through PRF Grant 2710-B.
Solid Solution of Aluminum Oxide in Rutile Titanium Dioxide1 by Richard A. Slepetys2and Philip A. Vaughan Department of Chemistry, Rutgers-The
State Uniuersitg, New Brunswick, New Jersey
(Received July $8, 1968)
The solubility of A1203in rutile Ti02 was determined in the temperature region of 1200-1426' from the change of the c unit-cell parameter. Solubility increased with equilibration temperature, ranging from 0.62% A1203 by weight at 1200' to 1.97% at 1426'. Corundum was the equilibrium solute phase below, and AlzTiOs above, 1240'. Solubility also increased with the addition of NblOS. The density of a rutile sample equilibrated at 1426' containing 1.6% by weight A1203was 4.199 g/cma, not significantly different from 4.202 g/cm3 measured for a corresponding pure rutile sample. These results are consistent with an interstitial in 1:2 proportion are placed on the respective rutile lattice sites, and solubility model in which AIS+ and 02the remaining A13+ occupies interstitial spaces. Treatment of the temperature dependence of the solubility data gave values for AH' of solution of 35 kcal/mol and 100 kcal/mol for A12Ti05 and AlzOs, respectively. The heat of formation of AlzTiOs computed from these results is -559 kcal/mol near the transition temperature. Introduction Several investigator^^-^ have studied the system AlzOa-TiOz. Their phase diagrams show one compound, A12TiO;, and eutectics with both component oxides. Lang, Fillmore, and XIaxwel16 also indicated a region of instability occurring in the range of about 750-1300' wherein aluminum titanate decomposed into the respective oxides. The solubility of aluminum oxide in rutile titanium dioxide had not been studied systematically. In her work on the preparation and structure determination of aluminum titanate, Hamelid noted that the solubility of aluminum oxide in rutile, if any, is very limited. F10rke7 reported an approximate value of 2 mol % at 1400". In the present investigation the extent of solid solubility of A1203 in rutile was measured in the region of the reported transition AbO3
+ Ti02
A12Ti05
The temperature of this transition was more firmly established, some thermochemical results were derived, and a solubility model was proposed.
Experimental Section A . Chemicals. The compounds used in this study were TiCL obtained from National Lead Co., Titanium Division, Baker Analyzed reagent AlCls 6H20, and Ciba anhydrous NbC16. E . Sample Preparation. Mixtures of oxides of titanium and aluminum were prepared by coprecipitation from a chloride solution. Measured amounts of analyzed stock solutions were mixed and diluted to contain about 0.07 g of Ti02/ml, the solution was heated to 55', neutralized with ammonium hydroxide to pH 6.5, filtered, washed, dried at 120', and ground in an agate +
(1) Based upon a dissertation submitted by R.A. Slepetys in partial fulfillment of the requirements for the Ph.D. degree in chemistry. (2) To whom correspondence should be addressed at National Lead Company, Titanium Division, P.O. Box 58, South Amboy, N. J. (3) H. yon Wartenberg and H. J. Reusch, Z . Anorg. A&. Chem., 207, l(1932). (4) E. N. Bunting, J . Res. Nat. Bur. Stand., 11, 725 (1933). (6) S. M. Lang, C. L. Fillmore, and L. H. Maxwell, ibid., 48, 298 (1952). (6) M. Hamelin, BUZZ. SOC.Chim. Fr., 1421 (1957). (7) 0. W. Florke, Ber. Deut. Keram. Ges., 38, 133 (1961).
Volume 73, Number 7 July 1969
