Aug. 20, 1957
AlQU.4TION O F C i s - D I C H L O R O - B I S -
[CONTRIBUTION FROM
THE
(ETHYLENEDIAMINE)-CHROMIUM(II1)
4285
W. A. NOYESLABORATORY OF CHEMISTRY,UNIVERSITY OF ILLINOIS]
The Kinetics of Aquation of cis-Dichloro-bis-(ethylenediamine)-chrornium(III) Ion1 B Y JOEL S E L B I N ' AND JOHN
c. BAILAR,JR.
RECEIVED MARCH30, 1957 The kinetics of acid hydrolysis of cis-[Cr end&] + has been studied a t PH 1.0 a t three temperatures and uuder various conditions of added electrolytes by a procedure which employs a chloride concentration cell t o measure chloride release. The pseudo first-order rate constant for the loss of one chloride is given by k, = 9.65 X 10" exp(-21,100/RT) set.-', and AS+ = -5.72 e.u. A value for the specific r a k constant for the loss of a second chloride has been estimated. Results of studies on t h e change in optical rotation of [-cis- [Cr ennC12] 7- are interpreted in terms of a n initial aquation without loss of optical activity followed by racemization of the chloroaquo product. Xnalogies with the corresponding cobalt complexes suggest t h a t complexes of the two metals follow similar mechanisms during acid hydrolysis. suggestion is made as t o why chromium(II1) complexes resemble cobalt(II1) complexes so closely. A\
Introduction
Relatively few kinetic experimentsg have ever been carried out on chromium(II1) complexes of A large amount of kinetic data has accumulated type (I) and only three have concerned complexes in the literature on simple substitution reactions in of type (11).~~~e'g octahedral complexes of the types: (I) Experiments by Adamson and Wilkinsgc on the or +1 and (11) [MA4YX]+1 Or +2 where 11 has been rates of aquation and exchange in the [Cr(NH3)6Co(III), Cr(II1) or Ir(II1) for type (I) and M has NCSIf2 and [ C O ( N H ~ ) ~ N C Sions ] + ~ led them to been Co(II1) and, in a few cases, Cr(II1) for type suggest that different mechanisms for substitution (11). The X represents a non-replaceable ligand may be operating for the two metals. Furtherand may be monodentate or polydentate, including more, the difference in electronic configuration exN- and,/or C-substituted polyamines. The X and isting between cobalt(II1) and chromium(III),10 Y, with X being the replaceable group, may be coupled with the invariably greater rates of aquaF-, C1-, Br-, I-, KCS-, NO2-, XOa-, N3-, tion found for chrornium(II1) complexes as comC03=,HC03-. OCHa-, CzH30z-, OH- or HzO. pared with the analogous cobalt(II1) complexes, The three types of substitution reactions which might lead one to suspect a difference in mechahave been studied are: (1) aquation, or replace- nism. Therefore, the purpose of the work rement by water in neutral or acid solution; (?) ported in this paper was to further compare the hydrolysis, or replacement by hydroxide ion in chemistry of cobalt(II1) and chromium(II1) combasic media; and ( 3 ) anation, or substitution of X plex ions. The compound chosen for study was by another anion.a cis- [Cr en&lz]Cl.H20. The reasons for this choice Although kinetic studies on complex ions were are: (1) the complex is readily prepared and reareported as early as 1911,4it has been only recently sonably stable for long periods of time (months) that the emphasis in the study of the kinetics has if kept in a desiccator in the dark; (2) the correbeen 011 the elucidation of the mechanism by which sponding cobalt(II1) complex has been thoroughly substitution reactions occur. studied5,"; (3) the complex undergoes aquation a t Practically all of the recent work on complexes a rate which can be followed conventionally; of types (I) and (11) has been concerned with the and (4) the complex is easily resolved into its opcomplexes of cobalt(II1). These have been studied tical isomers,. allowing studies of racemization. mainly by Basolo, Pearson and co-vvorkers,6Taube The aquation reaction, which has been studied, and co-workersJ6 and by Ingold, Nyholm and in 0.10 N acid solution, takes place in two steps Tobe' in aqueous solution, and by Brown, Ingold which may be written and Nyholm*in absolute methyl alcohol. (1) This article is based upon part of a dissertation submitted in partial fulfiltment of t h e tequirements for t h e P h . D . degree a t the University of Illinois, 1557. (2) Socony-hlobil Laboratolies Fellorr in Chemistry, 195-1957. (3) 1.' Basolo, Chem. Revs., 63, 459 i l Q 5 3 ) . ( 4 ) .4. B. Lamb and J. TV. Marden, THIS J O U R N A L , 83, 1873 (1911). ( 5 ) F. Basolo, J. G. Bergmann and R. G. Pearson, J . P k y s . Chew., 66, 22 (19.52); R. G. Pearson, C. R. Boston and F. Basolo, THIS J O U R K A L . 76, 3085 (1953); G. Basolo, B. D. Stone, J. G. Bergman and R . G. Pearson. i b i d . , 7 6 , 3079 (1951) ; R. G . Pearson. C. R . Boston and F. Basolo. J. Phys. Chem., 69,304 (1955); R. G. Pearson, K. E . Meeker and F. Basolo, THISJOURNAL, 7 8 , 709, 2673 (1956); F. Basolo, W. R. Matoush and R. G. Peatson, i b i d . , 7 8 , 4883 (1956); A. W. Adamson a n d F. Basolo, Acta Chem. Scand., 9, 1261 (1955). (6) A. C. Rutenberg a n d H . Taube, J. Chem. P h y s , 90, 825 (1952); H. Taube and F. -4.Posey, THIS J O U R N A L , 76, 1463 (1953); F. A. Posey and H. Taube, {bid., 79, 255 (1957). (7) C. K . Ingold, R. S. Nyholm and M. L. Tobe, J . Chem. Soc., 1691 (1956); R . S. Nyholm and M. L. Tobe, ibid., 1707 (1956); S. Asperger and C . K. Ingold, ibid., 2862 (1956). ( 8 ) D. D. Brown, C. K. Inxold and R. S. Nyholm, ibid., 2674 11953); D. D. Bromn and C. K. Ingold, ibid.. 2680 (1953): D. D. lkown and R. S. Nyholm, ibid., 2696 (1953).
cis-[Cr enzC1z] A
+
ki + H20 +
cis-[Cr en2CI(HnO)] i2
cis-[Cr eii2C1(H20)]r 2 B kz HnO
+ CI-
(1)
+ Cl-
(2)
+
k- z cis-[Cr en?(H20)n]+ 3 C
That the aquation proceeds stepwise has been shown in this study by both conductivity and kinetic measurements, although salts containing (9) H . Freundlich and R. Bartels, 2. physik. C h e m . , 101, 177 (1922); H. Freundlich and H . Pape, i b i d . , 86, 458 (1911); A. Adamson and R. G. Wilkins, THIS JOURNAL, 76, 3379 (1954); C. Postmus and E. L. King, J . Phrs. Chem., 69, 1216 (1955); D. M. Grant and R. E. Hamm, THISJ O U R N A L , 7 8 , 3006 ( 1 9 % ; M. Ikuta, H. G. McAdie and W. MacF. Smith, Can. J . Chem.. 34, 1361 (195G); R. E. H a m m , THISJOURNAL, 76, 609 (1953). (IO) H . Tauhe, Chetn. Revs., SO, 69 (1952) (11) J. P. Mathieu, Bd1. SOL chim. France, 3, 2121, 2156 (193ti).
4286
JOEL
SELBIN AND
JOHN
C. BAILAR, JR.
VOl. 79
the plates in a 1.O M hydrochloric acid-bath at a current of 25 ma. for 10 min. Electrodes prepared in this way were stable for several months and never had a n asymmetry potential greater than 0.2 mv. Between runs they were stored in 0.05 M hydrochloric acid in the dark. The cell was immersed in a constant temperature bath controlled to within =k0.05'. A Rubicon K-2 potentiometer and Leeds and Korthrup galvanometer were employed for the potential measurements. A stock solution of standard chloride ion concentration of 0.0100 M was prepared from NaC1 and either 0.100 M HC104 or 0.100 M "03. For runs p o t involving added electrolytes, the procedure was as follows: 50 ml. of standard chloride solution was placed in Vessel A. A weighed amount of complex was dissolved in 50.0 ml. of 0.100 M acid (either HC104 or "Os, but the same as was used t o prepare the standard solution) which previously had been brought to the proper temperature. This solution of the complex was placed in Vessel B and the solutions from both A and B were allowed to flow into the corresponding cell compartments with the pinch clamps closed over the Tygon tubing (see Fig. 1). The clamps were released after the compartments were filled ant1 Experimental junction was thereby made in the center bulb. All of this Preparation of Compounds.-cis- [Cr en2C1z]C1,H20 was required just under one minute from the time the solution of the complex was prepared. The rate of chloride producprepared according to the method described by Rollinson tion in the compartment containing the complex was then and Bailar.14 The recrystallization procedure was modified continuously and readily followed by observing the change as follows. The impure salt was dissolved quickly in water (4 ml. for each gram of salt) which had been acidified with in cell potential. The 0.100 Macid present in each solution served a twofold several drops of concentrated hydrochloric acid and the purpose. First, it prevented any reaction of hydroxyl ion solution was heated to about 65". The solution was filtered rapidly into ice-cold concentrated hydrochloric acid (1.5 from interfering with normal aquation of the complex, and second, it eliminated concern over a junction potential in ml. for each gram of salt used) and the filtrate was cooled t o the cell. -5" in an ice-salt mixture. T h e red-violet needles which T o be sure t h a t no chloride ion transfer occurred from one came out were collected on a sintered glass funnel, washed cell compartment to the other during the time of a normal well with alcohol and ether, and then dried in a vacuum desiccator. Equivalent conductance at zero time, A700 = run, standard chloride solutions of differing concentrations 103, which is in the correct range for a uni-univalent salt were allowed to run into the cell compartments and the pocontaining a complex cation.16 Anal. Calcd. for [Cr en2- tential was measured periodically. KOchange in potential Cl2]C1.H20: C , 16.19; H, 6.12; K, 18.88. Found: C, occurred over a period of time equal to that of a normal kinetic run. 16.14; H, 6.03; N, 18.40. The complex salt, cis-[Cr emThe runs which involved added electrolytes were carried Cl2]C1.H20,was resolved into its optical isomers by a modiout in the same way except t h a t a weighed amount of the fication of the method of Werner.16 T o a freshly prepared solution of [Cr enzC121Cl.HzO (12.0 g., 0.0404 mole) in 200 material whose effect upon the rate of aquation was to be determined was added to 50.0 ml. of the standard solutioti ml. water at 19' was added ammonium d-a-bromocamphorand to 50.0 mi. of the solution of the complex to produce the *-sulfonate (22.5 g., 0.0687 mole) with vigorous stirring. desired concentration. The solution was kept at 19". After several minutes the T h e run in "strong light" was made with a GO-watt bulb sulfonate salt separated. It was filtered and washed with approximately one foot from the cell. alcohol and ether (10.0 g., 89.552 diastereoisomer), and was Racemization Measurements.-Measurements of change then dissolved in concentrated HCI in the ratio of 4 ml./g. in optical rotation of h i s - [Cr en2C12]Cl.H20 were made in T o this solution was added two volumes of alcohol, which a 10 cm. jacketed polarimeter tube maintained at 21.9 precipitated fine, dark-red crystals. These crystals were filtered, washed well with alcohol and ether, and dried in a 0.1". The polarimeter used is a Schmidt and Haensch KO. 9143. Measurement: were made a t the sodium-D line or vacuum desiccator. The yield was 3.85 g. or 71.8%. The observed rotation for a 0.07570 solution was 4-0.235; approximately 5890 A . Solutions were 0.005 Jf (0.148c/b) [cY]*~D +314. Anal. Calcd. for I-& [Cr e n ~ C l ~ l C l ~ H ~in0 complex : and 0.100 X in HC10, or HXOa. At this conC, 16.19; H , 6.12; N, 18.88; Cr, 17.28. Found: C, centration, precision is initially poor due to the high absorb16.44; H, 6.12; N, 18.50; Cr, 17.54. All other reagents ancy of yellow light by the solution. This absorbancy dewere commercially available materials of reagent grade, creases fairly rapidly so that readings after about 20-2.5 and were used without further purification. minutes are more precise. The rotatory dispersion curve Kinetic Measurements.-Preliminary determinations of of this complex" indicates that the rotatory power is highcst tiic rate of loss of chloride from cis- [Cr end&]+ were made at a wave length of approximately 5490 A. However, a t by titrating the ionic chloride potentiometrically with silver this wave length practically all light is absorbed by the comnitrate after the reaction had been quenched by running plex solution, which is red-vialet in color. At the red eiid the aliquot being analyzed into an ice-cold 1: 1 ethanolof the spectrum (about 6000 A . ) very little light is absorbed water mixture. The results indicated t h a t the reaction is by the solution, but the rotatory power is sninll and the eye too fast to follow conveniently by a titration technique. is not very sensitive to changes in light intensity. IIcticc, Therefore, a more rapid and, at the same time, a more acthe sodium-D line was a compromise choice. curate means of following the loss of chloride, was sought. The method finally adopted was to employ a concentration Results and Discussion cell composed of two Xg-.-ZgCI electrodes. The cell design appears in Fig. 1. T h e Ag-AgC1 electrodes were prepared The iiieasured potential of the concentration cell by plating silver from a KAg(CN)z bath onto platinum was converted to a chloride concentration by means plates having a n area of 1 cm.2, followed by anodization of
ion B never have been isolated. As will be shown, reaction 1 proceeds a t a rate approximately 10 times that of reaction 2 , so that no serious interference is caused by reaction 2 in the study of reaction 1. The diaquo cation (C) is stable only in strongly acid solution and has been shown to retain the cis configuration.12 I t will revert to the cisdichloro cation only on standing for several day3 in a concentrated hydrochloric acid solution or upon evaporation of a concentrated hydrochloric acid ~ o l u t i o n . ' ~Therefore, under the conditions of the experiments reported here, it is reasonable to assume that equilibrium is not established either in reaction 1 or 2 and that in any case, an equilibrium in reaction 1 would lie far toward the right. Indeed, the experimental data allow an approximation of k z to be made.
+
of the relation
(12) P. Pfeiffer, Z. anorg. Chem., 68, 228 (1908). (13) P . Pfeiffer, ibid.. 66, 261 (1907). (14) C. L. Rollinson and John C. Bailar, Jr., "Inorganic Syntheses," Vol. 11, McGraw-Hill Book Co., Inc.. New York, N. Y . . 1946, p. 200. (15) H. J . Emeleus and J . S. Anderson, "Modern Aspects of Inorganic Chemistry," D. Van Xostrand Co., Inc.. N e w York, pi. Y . , 1955, p p 106, 108, 111. ( 1 6 ) A. Werner, Ber.. 44, 3132 (1911).
wherc E is the measured potential, n is the fixed standard chloride concentration, and x is the chloride concentration in the solution of the coni(17) J P M a t h l e u , Bfcil SOC c h t v i , 3 , 47b (1936)
hug. 20, 1957
4257
AQTJATIONOF C~~-~ICIII~ORO-~IS-(ETI-IYLENEDIAMINE)-CI-IRO~lIUhI (111)
it was possible by the method of successive approximations to obtain a value for k2 from known or measured values of a,kl, XI,x2 and t. The specific rate constant, kl, for reaction 3 is given in Table I for the conditions specified. I n most cases duplicate runs agreed within + 2y0, with + 3y0being the maximum discrepancy. The rate constant, kz, for reaction 4 is also included, but its value should be considered less accurate due to the method by which it was obtained. TABLE I RATESOF ACIDHYDROLYSIS OF cis-[M e n ~ C l ~ IN ] +
0.10 A 2 1 Temp., 'C
Complex
" 0 3
Ionic strength
k l , see.-! X 10'
kx, see.-' X 105
cis-[Cr enzcl?]+
Fig. 1.-Concentration
20.0 0.105 1.60 1.0 20.0 . l W 1.57 .. 25.0 .lo5 3.30 2.8 30.0 ,105 5.88 .. cis- [Co en~Clz]+ 25 .lo5 2.5* 25 .022d 1.2c 30 .022d 2.15' 25 0 .lo5 1.76e 0.05 M iTaC101. Value obtained by Pearson, et aZ.6d Values obtained by Mathieu." 0.012 iM HNOa and 0.01 Mcomplex used. e See Discussion.
cell for measuring rate of chloride ion production.
plex a t any time. The conditions a t early reaction times may be represented as ICr en~C121
+
(c
- Xl)
+ C1,- + H
ki 2 0 ---f
c
[Cr end21H201 + 2
+ C1.- + Clb-
XI
C
Several rate constant values for the cobalt complex are included and, as is seen, agreement among the several values is poor. Conductivity measurements were used by Mathieu" in obtaining the low value; spectrophotometric measurements were used by Pearson, et al.,6d in obtaining the high value; the intermediate value was obtained in our
(3)
x1
A t later times reaction (4) also must be considered
+ C1,- + clb- + H 2 0 c [Cr en?(H~O)zl + C1,- + clb- + (21,xz c k2
[Cr enzC1H2OI+* (Xi
--f
- x2)
x1
+3
x1
X2
-
(4)
-3.130-
I
I
I
I
+
The quantity measured experimentally is (c -3,050x1 x2) = (Cl-), but a t early times x2 0 and the value measured is (c XI). Hence, the concentration of [Cr enzClz]+ left a t any time is given by (c - XI) or (2c - [Cl-I), where c is the -2,030initial concentration. Plots of log (c - x1) vs. time were found to be linear up to a time representing about 60% completion of reaction 3. A sample plot appears in Fig. 2 . Thus, it is seen that reaction 3 is pseudo first order. The positive de- ,N viations from linearity were assumed to arise from the slower reaction indicated by eq. 4. Thus, a t times greater than about 40 minutes xz begins to 0, -become significant and therefore account must be G-2.690 taken of the fact that what is really plotted is (c - 3 x1 - xz). By extrapolating the straight line to longer times and taking the difference between the -2.570 value for the projected first-order decomposition and the actual value observed, a measure of x2 is obtained, ie., (c - XI - x.2) - (c - XI) = -x2. Using the standard equation18 for the series first-2.450 order reactions
-
+
+
c-2,810 ; ~
--Z.33UIX
which is (Xl
- x2)
k a1 kn
- ki
- e - w)
(18)A. A, Frost and R, G , Pearson, zsKineticsand Mechanism,so John Wiley and Sons, Inc., K'ew York, h'. Y.,1953.p. 153.
0
I
I
10
20
Fig. a.-First-order
30 40 50 60 70 80 90 Time, min. rate plot for acid hydrolysis of cis[ CrenzCls] +.
concentration cell. The latter measurements were not as precise as was expected, due to interaction between the cobalt complex and the Ag-AgC1 electrode, preventing the rapid establishment of equilibrium conditions between solution and electrode. Thus, meaningful readings could not be taken for a t least ten minutes after the solution of the complex was placed in the cell, and experimental scattering was increased. The chromium complex did not affect the electrodes in this manner. No further study of this phenomenon was made. However, as expected, the chromium complex aquates faster than the corresponding cobalt complex. The thermodynamic data for reaction 3, which were obtained from a linear drrhenius plot of log k zls. 1/T are: AE* = 21.1 kcal., AH' = 20.5 kcal., AF12g80 = 22.2 kcal., A S * = -5.72 cal./deg. and log Z = 11-98. Corresponding values for the acid hydrolysis of the cobalt complex, cis-[Co en? C12]", are": AE* = 22.25 kcal., AH+ = 21.6 kcal., and log Z = 14.02. The energy required to form the activated complex is therefore similar for both metals, but is slightly lower in the case of chromium, as is expected. Table I1 records the rate of acid hydrolysis a t 23.0" and a pH = 1.O (obtained with HC104) under
to prevent attack by the negative end of a water dipole which is necessary for the Sii2 mechanism. However, the aquation of czs- jCr en&lz]+ is not as greatly affected by the presence of sulfate ion as arc the aquations of [Cr(NH3)&1]+?9f and [Co(XHJ5C1]+2.19 This result was expected since the latter ions are doubly charged whereas the dichloro cation is singly charged. The slight variation in rate observed with a change in initial concentration is not a t present understood and requires further study. A study of the change in optical rotation of 1-cis[Cr enzClz]+has been made and the data may be interpreted in a manner similar to that used to treat data obtained with the corresponding cobalt complex.20 Thus, in a plot of log a: zrs. t , an initial straight portion is joined to a final straight portion by a curved line. A s 1IathieuZohas shown for the cobalt salt, the initial straight line yields a k value close to the acid hydrolysis constant, kl, and the final straight line yields a specific rate constant for racemization of the species being formed by the first process, which in this case is I-cis-[Cr end21 HzO]+?. If k2 (Table I) is compared with k , (Table 111),it is seen that the optically active chloroaquo cation racemizes a t a faster rate than i t loses chloride. This is similar t o what has been observedz0in the case of the analogous cobalt complex and indicates TABLE I1 that a different mechanism is operating to produce RATESOF .4CID HYDROLYSIS O F CiS-[Cr enoclz]' UKDER racemization than is operating to produce further \ r CONDITIONS ~ AT ~ 25.0' ~ASD p H~ 1.0(HClOr) ~ ~ aquation. -4 similar treatment has been given for Concn. of the change in rotation of d-cis-[Co en2 Fz]'.5f Added electrolyte or complex, Ionic kl, set.-' other condition strength x 104 x0. .M x 103 Table I11 records the rates of acid hydrolysis, ................ 0.105 1 5.00 3.30 mutarotation and racemization for the above com60 w a t t bulb 1 foot 2 5.00 pounds. It appears that the initial change in rotaaway 3.27 .lo5 tion is due mainly to the formation of the optically Optically active form 3.38 ,105 3 5.00 active chloro-aquo compound, which then under................ 3.22 ,106 4 6.00 goes racemization. ................ 3.47 ,104 5 4.00 It would seem from this work that the ions cis................ 3.58 ,102 6 2.50 [Co en4&]+ and cis-[Cr en2C12]+are not very dif0.05 A 1 Mg( C104)? 3.15 ,255 7 5.00 ferent in their behavior, at least toward acid hydrol.05 M Xi( ClO4)z 3.20 8 5.00 ,255 ysis and optical change. This observation sug3.63 .05 ,If Sac1 9 5.00 ,115 gests that the two metal complexes follow a similar ,010M A-asSOi 3.30 .135 10 5.00 mechanism during the course of acid hydrolysis. ,025 31 SazS04 ,180 3.28 11 5.00 Although the purpose of this work was not to at3.57 ,050 .If SazSO4 12 5.00 ,255 tempt to decide upon a mechanism, the results ,075 -11 Saos04 3.68 .5.00 13 ,330 would seem to favor for the chromium complex the 100 -If XazSO4 4.03 ,405 14 5.00 Sx1 mechanism suggested by Basolo, Pearson, various conditions of concentration and added ct d . , j for the acid hydrolysis of cobalt complexes. If the assumption is true that both cobalt(II1) clectrolyte. Several conclusions can immediately be drawn. The presence or absence of light (no. and chromium(II1) are using 3d?4slp3hybrid orbi-3) has no apparent effect on the reaction studied, tals in their complex compounds, then the remainas is apparently the case for some cobalt com- ing three 3d orbitals are filled in cobalt and halfplexes.7a Variation in ionic strength (nos. 7 , 8) ap- filled in chromium. It may well be that the strong pears to have no effect on the aquation reaction; similarity found between cobalt(II1) and chrothis is in line with Mathieu's results with the co- mium(II1) complexes is related to this fact. 'The balt complex.l1 However, chloride ion and sulfate similarity in reactivity between elements having ion (nos. 9-14) in concentrations 10-fold or greater the same orbital set either empty, half-filled, or over complex concentration appear to accelerate filled has been noted before in both the transition the aquation reaction. This can be accounted for and inner transition elements. Such orbital arby assuming ion-pair formation between anion and rangements seem always to lend to greater stability dichloro cation. If this is indeed the case, it con- than would otherwise be predicted. Thus it is stitutes additional support for an S s l or dissocia- suggested here that the filled orbitals in the cobalt tion type mechanism. Thus, an ion pair [Cr en? complexes and the half-filled orbitals in the chroC12]+-S04= would induce an easier breaking of (19) F J Garrick, Trans F a r a d a y S O L , 34, 7088 (1938) the Cr-C1 bond and at the same time would tend ( 2 0 ) 1 P V a t h i e t i , F?r/l TOC chiin 4 , 087 (1037)
TABLE I11 RATESO F ACIDHYDROLYSIS, MUTAROTATION AND RACEMIZATION GIVEKI N k lnitial compd.
I-&-
4289
THERMODYNAMIC PROPERTIES OF PYRIDINE
Aug. 20, 1957
(Acid)
[Co en?Cl?]+a
d-cis- [Co enlFn] +b 2 4 s - [Cr ennCl?]
Values from Mathieu.?O
*
ki, acid
Temp., OC.
0.012 30 ,001 25 .1 25 .1 22 .1 25 Values from Basolo, et
hydrolysis
2.15 x 10-4 6.0 X 5 x 10-5 2 . 3 0 x 10-4 3 . 3 0 x 10-4
FROM THE
kc, racemization
5.37 x 8.2 X 10.3 x 2.28 x 3.30 x
6.13 X 3 x 6.5 X 1.48 x 2.0 x
10-4
10-5 10-4 10-4
10-6 10-6 10-4 10-4
d.5'
mium complexes are in part responsible for (1) the general non-labile characteristics found for cobalt and chromium complexes as compared with other first row transition metal complexes, ( 2 ) the strong similarity of cobalt and chromium complexes in
[CONTRIBUTIOS K O . 62
(SEC.-')
kn, mutarotation
ease of substitution-type reactions, and (3) the ease with which Cr(I1) goes to Cr(II1) in systems which can form d2sp3 hybrid-orbital complexes with Cr(II1). ERBASA,ILLINOIS
THERMODYSAMICS LABORATORY, PETROLEUM EXPERIXEYT S T A T I O N , BUREAU O F MIXES, U. s. DEPARTMEST O F THE I N T E R I O R ]
Pyridine : Experimental and Calculated Chemical Thermodynamic Properties between 0 and 1500OK.; a Revised Vibrational Assignment1 BY J. P. MCCULLOUGH, D. R. DOUSLIN, J. F. MESSERLY, I. A. HOSSENLOPP, T. C. KINCHELOE A N D GUY WADDINGTON RECEIVEDAPRIL 22, 1957 From studies by low temperature calorimetry, comparative ebulliometry, flow calorimetry and combustion calorimetry, values were obtained for the entropy, heat capacity and heat of formation of pyridine in the ideal gaseous state. The calorimetric values of entropy and heat capacity served as guides in revising the vibrational assignments for pyridine and several deuteropyridines. T h e revised assignment for pyridine and molecular structure d a t a were used in computing values of the thermodynamic functions a t selected temperatures between 0 and 1500°K. Values of the heat, free energy and equilibrium constant of formation were calculated a t the same selected temperatures. T h e experimental investigations provided t h e following information: values of heat capacity for the solid (13-224'K.), the liquid [ C s a t d = 33.633 - 9.1446 X 10-ZT 3.8478 X 10-4T2 - 3.3430 X 10-'T3, cal. deg.-'mole-' (240-347'K.)], and t h e v a p o r [C," = -8.262 10.608 X 1O-*T 5.4662 X 1 0 - 5 P , cal. deg.-' mole-' (374-50O0K.)] ; the heat of fusion [1978.6 cal. mole-'] a t the triple point [231.4g°K.]; t h e entropy of the liquid a t 298.16"K. [42.52 cal. deg.-' mole-']; the vapor pressure [log," p(mm.) = 7.04162 - 1374.103/ (t 4- 215.014), (67-153")]; the heat of vaporization [ A H v = 13077 - 10.201 T - 4.7941 X 10-3T2, cal. mole-' (346388'K.)]; the second virial coefficient in the equation of state, P V = R T (1 f B / V ) , [ B = -15-84.5 exp ( 9 5 0 / T ) , cc. mole-' (346-500°K.)] ; a n d the standard heat of formation of t h e liquid a t 298.16'K. (23.89 kcal. mole-'].
+
Thermodynamic investigations of organic nitrogen compounds were begun recently in this Laboratory as part of American Petroleum Institute Research Project 5 2 . The objective is to provide comprehensive thermodynamic data for nitrogen compounds that are important in petroleum technology. This paper is the first of a projected series on cyclic nitrogen compounds. Results are reported for pyridine, a compound important in theoretical considerations concer=iing molecular structure as well as in practical considerations as an industrial chemical. Some thermodynamic data for pyridine have been reported previously,2 but many of the basic data are not of the accuracy that can be obtained by present-day methods. (1) (a) This investigation was performed as p a r t of American Petroleum Institute Research Project 52 on "Nitrogen Constituents of Petroleum," which is conducted a t t h e University of Kansas in Lawrence, Kansas, and a t the Bureau of Mines Experiment Stations in Laramie, Wyoming, and Bartlesville, Oklahoma. (b) Presented a t 130th American Chemical Society Meeting, Atlantic City, AT, J., September, 1956. (2) (a, C. H. Kline and J. Turkevich, J . Chem. Phrs., 12,300 (1944); (b) D. P. Biddiscombe, E. A. Coulson, R . Handley and E. F. G. Herington, J . Chem. SOL., 1957 (1954); (c) E . F. G. Herington and J. F. Martin, Trans. F a r a d a y Soc., 49, 151 (1953) [and references cited JOURNAL, there]; (d) G. S. Parks, S. S. Todd a n d W. A. Moore, THIS 68, 398 (1936); (e) J. N. Pearce a n d H. >I. Bakke, Proc. I o w a A c a d . Sci., 38B, 450 (1929).
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The experimental part of this investigation consisted of measurements of thermodynamic properties by the methods of low temperature calorimetry, comparative ebulliometry, flow calorimetry and combustion calorimetry. Descriptions of the experiments and detailed results are given in the Experimental part of the paper. For convenience, the most important experimental results-values of the entropy, heat capacity and heat of formation for the ideal gaseous state-are collected in Table I. The entropy and heat capacity data served as TABLE I OBSERVED AXD CALCULATED MOLAL THERMODYNAMIC PROPERTIES OF PYRIDINE IN THE IDEALGASEOUS STATE T,
OK.
346.65 366.11 388.40
Entropy, S o , cal. deg. -1 Obsd. Calcd.
70.65 71.86 73.27
T,
H e a t capacity, CpO, cal. deg. -1 Obsd. Calcd.
OK.
70.65 71.89 73.30
374.20 23.79 23.79 397.20 25.23 25.24 420.20 26.65 26.65 450.20 28.43 28.40 500.20 31.12 31.21 Heat of formation, AHf0?9S16 (obsd.) = 33.50 kcal.
guides in revising the vibrational assignment for pyridine and as checks on thermodynamic properties calculated by the methods of statistical
