NOTES
July, 1957 solution. Nickel chloride solutions were standardized gravimetrically by quantitatively precipitating nickel as nickel dimethylglyoxime. The sodium hydroxide solutions were standardized by titration against primary standard potassium acid phthalate using phenolphthalein as an indicator. The mixture calorimeter used for these measurements ronsisted of a glass vessel open a t both ends, having a capacity of 250 ml. This glass vessel fit down into a one-liter Dejvar flask. A Beckman thermometer was placed through the neck of the glass vessel, the bottom of which was sealed with a wax disc. When it was desired to mix the two solutions the inner glass vessel was raised, causing the wax disc to be knocked out by the Beckman thermometer. The solutions were stirred by continuously raising and lowering the glass vessel. The calorimeter was standardized by measuring the neutralization of NaOH with HC1. From the average of ten measurements of the heat of neutralization of this reaction a calorimeter constant of 60.4 cal. per deg. was obtained. The calibration was checked by measuring the reaction of BaCls with HzSOa. The values obtained for the heat of this reaction agreed with literature values to within less than 1%. The heat of neutralization of tetren.5HCl with NaOH was measured by mixing 200 ml. of 0.4709 N tetren.SHC1 with 225 ml. of 0.4694 N NaOH. The results obtained are given in Table I together with the data obtained from the constants measured by the method of Schwarzenbach.s
1007
obtained by thermochemical measurements a t 25” was -20.60 kcal./mole. This value agrees well with the values calculated from the slope of log K 21s.1 / T using the basicity constants reported previously. The calorimetrically measured heat of iieutralization does not, however, agree with the values obtained by Schaafsma6 from acid constants measured by the method of Bjerrum. She reports the heat of neutralization between 25 and 35” as -35.G9 kcal./mole. These comparisons illustrate again that the variation in the constants with temperature obtained by these methods are not reliable criteria for the calculation of thermodynamic data. The average value obtained for the heat of the reaction (I) Was - 1l.31 kcal./mole. The heat of reacNil+
+ 2C1- + tetren
(NitetrenH20)2+
+ 2C1- ( I )
tion (11) used for icomparison of the thermodynamic Ni2+
+ 2C104- + tetren
( NitetrenH20)2
+ 2C104- (11)
quantities reportled in Table 11, was calculated
from the heat of reaction I, and standard heats of TABLE I formation for the Ni2+, C1- and Clod- ions using THERMODYNAMIC DATAFOR THE HEATOF NEUTRALIZATION OF TETRAETHYLENEPENTAMINE.5Hcl
AH, kcal./mole
1 Thermochemical measurements 2 Values from slope log K us. 1/T AP,
t, oc.
-20.60 -18.15 -21.41
25 15-25 25-35
kcal./mole
3 From constants by method of Schwarzenhach
47.13 48.15
25 35
AS, cal./mole/deg.
Thermochemical measurements From data in 2 and 3 above
-231 - 228 - 229
25 25 35
The heat of reaction NPf 2C1tetrcn _J (NitetrenH20)*+ 2CIwas measured by mixing 200 ml. of 0.2341 If NiClz with 200 ml. of 0.245 M tetren. The value obtained is given in Table I1 together with the values calculated from the variation of the formation constants with temperature as measured by the method of S c h ~ a r z e n b a c h . ~
+
+
+
TABLE I1 THERMODYNAMIC DATAFOR Nif2
THE
REACTION
+ 2C104- + tetren = Piitetren2’ + 2C104AH,
kcal./mole
Thermochemical measurements
From formation constants AF, kcal./mole
From formation constants
- 10.43
t, o c .
-10.26
25 15-25
-11.11
25-35
-23.25 -23.89 -24 53
15 25 35
-45.1 -45.1 -44.3 -43.5
25 15 25 35
i
AS, cal./mole/deg.
Thermochemical measurements From formation constants
i
Discussion The average value for the heat of rieutrtllizatioii
literature values.8 The calculation yielded a value of - 10.43 kcal./rnole. The value obtahed from the formation constants measured by the method of Schwarzenbach is - 10.68 kcal./mole between 15 and 25”. Acknowledgment.-The authors wish to thank the Office of Ordnance Research, U. S. Army, for support of this and continuing investigations. (8) “Selected Values of Chemical Thermodynamic Properties,” IJational Bureau of Startdards, Washington, D.C., 1947-50.
-
THE TRIBROMIDE EQUILIBRIUM I N AQUEOUS ACETIC ACID BY T. W. NAKAGAWA, L. J. ANDREWA N D R. M. KEEFER Contribution from the Department of Chemistru, University of California, Davis, California Received February 18, 1867
To provide information necessary for the quantitative interpretation of the inhibition of aromatic bromination reactions by bromide ion in acetic acid1v2and in aqu’eousacetic acid3 solutions, which occurs because of tribromide formation, the tribromide equilibrium ‘constant has been measured using a variety of acetic acid-water mixtures as solvent. Both hydrogen bromide and sodium bromide have been used as sources of bromide ion. The results of these measurements are sufficiently interesting to deserve special comment. Although some data of the kind reported here are already r e ~ o r d e d , 2 ~ ~ ~ ~ 4 (1) P. W. Robertson, J . Chem. Soc., 1267 (1954). (2) R. M. Heefer, A. Ottenberg and L. J. Andrews, J . A m . Chem. SOC.,78, 255 (1956). (3) See for example (a) A. E. Bradfield, G. I. Davies and E. Long, J. Chem. Soc., 1389 (1949): (b) R. M. Keefer and L. J. Andrews, J . A m . Chem. Soc., 78, 3637 (1956); ( c ) E. Berliner, Abstracts of Papers, American Chemical Society Meeting, Atlantic City, N. ,J., Sept., 1956, p. 56-0. (4) (a) W. J. Jones, J . Chem. Soc., 392 (1011); (b) K. Noaaki and R. A. Ogg, J . A m . Chem. Soc., 64, 697 (1942); (0) E. Grovenstein, Jr., and U. V. Heiiderson, Jr., ibid., 7 8 , BGO (1933); (d) L. J. A n d r e w and R. hI. Keefcr, ibid., 78, 4549 (195G).
1008 200
NOTES
r
I
I
I
I
I
I
'
I
I
I
Ri
I I
I
Vol. 61
The total concentrations of sodium bromide or hydrogen bromide in the solutions were used as (Br-) values. The term E B ~ Trepresents the ap arent extinction coefficient of bromine, both free and in t i e form of tribromide Other details of the method already have been described.a*;bi'd Table I presents typical results obtained by graphical analysis of the data. The measured values of € 1 3 ~ ~ -at any given wave length did not vary greatIy with changes in the bromide source, nor with the composition of the solvent. The K values as a function of solvent composition are presented graphically in Fig. 1.
TABLE I
TYPICAL VALUESOF x,
mu
340 350 360 370
eBra
E B ~ Z6, ~ ~ AND 3 -
tBrI-
50% HOAc, NaBr" 57 2360 78 1440 95 1040 120 880
K (25 i 0.2")
1. mole-'
28.1 28.4 27.6 26.4
tBr8-
1. mole-
80% HOAc, HBr 2350 99 1430 103 1020 98
...
...
75% HOAc, NaBr' 100% HOAc, NaBr" 54.3 2360 . 2380 34 340 53.4 1460 . 1470 61 350 51.9 1010 1030 94 360 51.3 860 860 370 127 a The total salt concentration of sodium bromide solutions was adjusted to 0.1 M by addition of sodium acetate. b These values of EB~,- were obtained by direct measurement of a solution which was 0.81 M in sodium bromide and 1.49 X lo-' M in bromine.
.. .. . .. . ..
20
40 60 80 100 Vol. % acetic acid. Fig. 1.-The tribromide equilibrium constant in aqueous acetic acid a t 25.0 i 0.2". The filled circles are based on dat,a taken using hydrogen bromide as the bromide source. The open circles are based on data taken using sodium bromide as the bromide source with total electrolyte concentration adjusted to 0.1 M by the addition of sodium acetate. The point at 0% acetic acid is taken from the work of R. 0. Griffith, A. McKeown and A. G. Winn, Trans. Faraday SOC.,28, 101 (1932), and H. A. Liebhafsky, J. Am. Chem. SOC., 56, 1500 (1934).
they have been coIIected under a diversity of experimental conditions and are not easily correlated. Experimental Materials.-The purification of the acetic acid, and the drying of salts have been described elsewhere.3b In preparing solutions of hydrogen bromide in pure acetic acid the gas was generated by dropping bromine on tetralin in a system protected from moisture. Solutions thus prepared were analyzed gravimetrically. Solutions of hydrogen bromide in aqueous acetic acid were pre ared from 48% aqueous hydrobromic acid which was anagsed volumetrically. The Equilibrium Constants.-These were evaluated spectrometrically, a procedure which is readily applicable because of the intense absorption of tribromides in the near ultraviolet region. In general five solutions of the same solvent composition were prepared in which the concentration of hydrogen bromide or sodium bromide varied from 0.01-0.1 M and the concentration of bromine varied from 3 X lo-' to 6 X lo-* M . Optical densities of these were measured a t several wave lengths in the 340-370 mp region using 1 cm. absorption cells and the solvent as a blank. In cases in which sodium bromide was used as the halide source, sufficient sodium acetate was added to fix the total salt concentration at 0.1 M . The molecular extinction coefficients, EB-, of free bromine in each solvent mixture were established by separate measurement. The equilibrium constants for the interaction (equation 1) were evaluated
K = (Bra-)/(Br-)(Brs)
(1) graphically from the linear plots of values of I/(EB~*T ~ B J obtained a t each wave len th against corresponding values of l/(Br-) (see equation 20% )
1/(EBrrT
-
BBR)
=
1/(BI'-)(Ic)(eBrs-
- €BP%)+ l/(BBn- - eBm)
(2)
I
K,
R,
Discussion As can be seen from Fig. 1 the K values are essentially independent of the bromide source until the acetic acid content of the solvent exceeds 90%. The relatively close agreement of the equilibrium constants obtained using sodium bromide and hydrogen bromide as halide sources in 90% acetic acid is not unexpected, since the two bromides function with comparable effectiveness in inhibiting mesitylene bromination in this rnedi~m.~bThe increases in magnitude of the K values as the acetic acid content of the solutions increases to 90% undoubtedly reflect the diminishing capacity of the solvent, as its strength as a Lewis base is decreased, to compete with bromide ion in interacting with bromine. Although the reported constant for the interaction of hydrogen bromide with bromine in pure acetic acid is of limited accuracy,b it is clear that in the anhydrous medium the extent of interaction of both hydrogen bromide and sodium bromide with bromine is much less than in 90-95% acetic acid. It is certain that in pure acetic acid the degree of dissociation to ions of both sodium bromide and hydrogen bromide is very small,6and the products of interaction with bromine must also be un-ionized (NaBr8or KBr3). However, hydrogen bromide, at least, is very likely appreciably dissociated to ion pairs in pure acetic acid.6 The capacity t o interact with bromine may be restricted to the ion pair forms of the two bromides, the abundance of which (5) The experimental data, when plotted according t o equation 2, did not oonforrn too well to linearity, possibly because dilute solution laws do not hold in anhydrous acetic acid. The K value plotted in Fig. 1 ie an average of values obtained by calculating the tribromide concentration of the solutions using measured optical densitiea and average values of C B ~ Z - taken from Tehle I. (6) I. M. Kolthoff and S. Bruckenstein, J . Am. Chem. Roc., 78, 1 (1956).
1
NOTES
July, 1957 must diminish significantly as the water content of the solvent falls to zero. The fact that the ultraviolet spectra of the tribromide species are relatively insensitive t o changes in solvent composition (see Table I) suggests that the tribromides formed in the anhydrous medium retain sufficient ionic character to display the spectrum typical of tribromide ion. The heat of reaction of bromine and sodium bromide in 75% acetic acid (AH" = -3.4 kcal./mole) has been calculated from the equilibrium constant for the reaction at 25 f 0.2" and a t 2.0 f 0.2". The K value for the reaction a t the lower temperature (100 l./mole) was determined by the same spectrophotometric procedure used in making measurements at the higher temperature.?
1009
3:
281-
'\
0
O
4r4 1 . 3 ~ H
I
1.2 e v
Y
1-14 1.o
(7) The values of t ~ ~ , -obtained in the measurements s t 2" were in good agreement with those obtained at 25O.
I 0.9
THE EFFECT OF TOLUIDINES AND XYLIDINE ON MALONIC ACID BY LOUISWATTSCLARE Contribution /Tom fhe Department of Chemistr , Saint Joseph CoZolleee, Emmitsburg, Marylana! Received March I d , iD67
The literature contains kinetic data on the decarboxylation of malonic acid in the following aromatic amines : aniline, quinoline, pyridine, 2-methylpyridine, 3-methylpyridine and 4-methylpyridine.'-8 The kinetics of this reaction have been investigated in this Laboratory in four additional aromatic amines, namely, 2-methylaniline, 3-methylaniline, 4-methylaniline and 2,6-dimethylaniliiie. The results obtained, which are reported herein, are consistent with the expected effects of ortho, meta and para substituents on the formation of the activated complex. Experimental Reagents.-( 1) Malonic acid, Analytical Reagent Grade, 100.0% Assay; (2) ortho-toluidine, Reagent Grade, b.p. 83-85' (15 mm.); (3) m-toluidine, Reagent Grade, b.p. 92-93' (15 mm.); (4) p-toluidine, Reagent Grade, m.p. 43-44'; (5) 2,6-dimethylaniline, Reagent Grade, b.p. 213-214' (760mm.). Apparatus and Technique.-The kinetic experiments were conducted in a constant temperature oil-bath (f0.1') by measuring the volume of COZevolved at constant pressure, as described in a previous paper in this ~ e r i e s . ~I n each experiment a 0.1857-g. sample of malonic acid (the amount required to produce 40.0 ml. of COZa t STP on complete reaction) was introduced in the usual manner in the reaction flask (a 100-ml. 3-neck standard taper Pyrex Brand flask containing 50 ml. of solvent saturated with dry COz gas). Temperature measurements were made using a thermometer calibrated by the Bureau of Standards, graduated in tenths of a degree, to which appropriate stem corrections and other necessary corrections were applied.
4
0.8
0
I1
0
50
100 150 200 250 300 350 Seconds. Fig. 1.-Experimental data, decomposition of 0.1857 g. of malonic acid in 50 ml. of m-toluidine a t 136.6' (cor.): I, ml. Cot at STP: 11,log ( a - 2).
of the reagent before decarboxylation could become effective. When the corrected volume of C 0 2 was plotted against time and graphs made of log (a - x) vs. t (where a is the maximum theoretical volume of C02 and x is the volume evolved at time t) from representative points on the smoothed experimental plots, straight lines resulted for the first 75% of the reaction. The primary decarboxylation reaction was therefore first order, and the rate at each temperature was calculated from the slope of the line. The data thus obtained are shown in Table I, and results for a typical run are shown in Fig. 1. TABLE I DECOMPOSITION OF MALONICACID IN TOLUIDINES AND XYLIDINE. RATE CONSTANTS AT VARIOUS TEMPERATURES Solvent
+Toluidine In-Toluidine p-Toluidine
Temp., O C .
k X 105 (seo.-l)
120.5 127,5 132.8 117.05 126.56 136.6 117.9 124.8 126.5 133.4 117.5 126.5 133.3
103 178.5 271 53 125 297 87 164 196 367 74.5 147.5 233
Results and Discussion 2,6-Dimethylaniline I n aniline solution each sample of malonic acid gave a 100% yield of C02 at the end of the experiment.2 In toluidines and xylidine the yield of COz The experimental data in Table I yielded straight was never more than 90%. Evidently in the latter lines when log k was plotted against 1/T (see Fig. solvents a slow secondary reaction consumed some 2). From the slopes of the lines in Fig. 2 activa(1) G. Fraenkel, R. L. Belford and P. E. Yankwich, J . Am. Chem. tion energies and frequency factors from the ArSoc., 76, 15 (1954). rhenius equation were obtained. The thermody(2) L. W. Clark, THISJOURNAL, 60, 1340 (1956). namic quantities in the Eyring equation were also (3) L. W.Clark, ibid., 60, 1583 (1986). calculated. These data are shown in Table 11. (4) L. W. Clark, ibid., 60, 1150 (1986).
