3340
NOTES
Table I : Equilibrium Constants for the Dimerization of 2-Pyridone and 2-Thiopyridonea r -
Temp, O C
16.2 20.3 26.8 30.3 31.9 35.3 37.5 41.6 42.5 45.9
2-Pyridone
10.43
x
Kd, M-lp-2-Thiopyridone
103
6.27
x
103
4.49
x
103
2.71
x
103
2.80
x
103
5.35 3.93
x x
103 103
3.22
x
103
2.70
x
103
2.30
x
103
a The probable errors in the slope and intercept calculated a t the 90% confidence level give standard errors of approximately 10% in the values of Kd.
Table I1 : Thermodynamic Functions for Dimerization of Thiolactams and Lactams
Name
2-Pyridone" 2-Thiopyridone" Thiobutyrolactam' Butyrolactamd Thiovalerolactam' Valerolactam" Thiocaprolactam" Caprolactam'
AHo, koal/mol
-8.8 -7.0 -5.8 -7.0 -3.1 -10.3 -3.9 -5.5
AG026, kosl/mol
f 0.4b -5.25 zt 0.08 -4.94 =k 0.06 -3.34f0.06 It 0 . 7 rt 0 . 4 -3.32 -3.62 f 0 . 0 6 f0 . 4 f 1 . 0 -3.34 -3.52 f 0.06 f0.5 It 0 . 3 -2.77 f 0.3b
ASa, eu
-12.0 -6.7 -8.3 -12.3 1.7 -23.4 -1.6 -9.2
' This work. Errors calculated using probable errors in slope and intercept. ' Reference 2. Reference 7. e ReferReference 3. ence 8.
'
thioamide group interacts with the double bonds of the ring. The AHo of dimerization for the thio compound is about 80% that of the oxygen compound, which is consistent with the ratios found for two out of three saturated systems previously investigated. Valerolactam is again the exception, with a much higher AH' than its thio analog. That the differences in AH" are so small can be explained on the basis of the increased acidity of the NH group in the thioainides partially compensating for the lower basicity of the C=S group.2 Thioanides would be more acidic, owing to the smaller resonance interaction of sulfur compared with oxygen and a consequent higher polarity. Finally, the values of Kd calculated at 25" are 7100 M-l for 2-pyridone and 4200 M-' for 2-thiopyridone. Unlike the saturated thiolactams and lactams, the AH" values for these compounds are large enough to dominate the entropy terms which had caused the saturated oxygen and sulfur compounds to have very similar Kd values. The values of Kd reported by Krackov, et uLJ1 show larger differences than those given here. This could be due t o the different solvents and higher concentrations used in their work. Acknowledgment. We wish to thank the National Science Foundation for the support of William Reineke as a participant in the Research Participation for High School Teachers Program, Grant No. GW-1702.
Kinetics of the Gas-Phase Reactions of Diborane with Methylphosphines and Trimethylamine
by H. Brumberger and W. H. Smith that both of these compounds form relatively strong H bonds. I n fact, the AH" value for 2-pyridone is about the same as some of the values observed for benzoic acids.4 Bellamy and Rogash came to a similar conclusion, basing it upon the large shifts of the KH stretching frequencies observed for the dimers of these compounds.6 A comparison of the data for the unsaturated and saturated compounds shows that the H bond is strongest for 2-pyridone in the oxygen series and 2-thiopyridone in the sulfur compounds. The one piece of evidence contrary t o this is the very high AH" (10 kcal) reported for valerolactam.8 However, this value is open to question, as the experimental techniques used (measuring the dimer peak instead of the monomer peak) can lead to large errors in That the unsaturated compounds are the more strongly H bonded can be attributed to the higher acidity of the NH groups and the higher basicity of the C=S and C=C groups. The higher acidity and basicity can be caused by contributions of resonance forms in which the amide or T h e Journal of Physical Chemistry
Department of Chemistry, Syracuse University, Syracuse, N e w York 15210 (Received A p r i l 8, 1968)
The reaction of phosphine with diborane in the gas phase a t 0" is relatively slow (kl % 2.3 cc mol-' sec-l) and can be accounted for by the mechanism
BzHe
+ PH3
ki
BHJ'H3(g) k- 1
+ BH3
The activation energy E1 = 11.4 f 2.0 kcal/mol;l the solid product is reversibly dissociated and has a dissociation pressure of 200 mm a t Oo.2 Observations on t,he reactions of diborane with (1) H. Brumberger and R. A. Marcus, J . Chem. Phys., 24, 741
(1956). (2) E. L. Gamble and P. Gilmont, J . Amer. Chem. Soc., 62, 717 (1940).
3341
NOTES methyl- and dimethylphosphine (MP and DMP) indicated that these were very fast, comparable in velocity to the reactions of diborane with amines. Some kinetic flow-reactor studies of the competitive reactions of JIP and D M P with diborane and of DMP and trimethylamine (TMA) with diborane are reported here. All materials were purified by degassing and distillation where nece~sary.~Vacuum and flow reactor systems were in all essentials similar to those described by Daen and Marcus14as were the techniques used to measure reaction rates, exchange reactions, etc. It was found by infrared spectroscopy of various reaction mixtures that exchange reactions between adducts and reactants were negligible, and that the pumping system used to collect effluent reactor gases for analysis did not significantly fractionate these gases. Infrared Beer's law curves established for pure components at appropriate absorption peaks were used to determine the composition of reaction mixtures to better than 5% for each component. All measureiments were made under essentially isothermal steady-state flow conditions. An excess of competitors relative to diborane was always present. Flow rates were obtained by dividing the pressure of the reactor effluent gas collected in a standard volume during a run by the elapsed collection time. Such flow rates (mm miiz-') were measured for the competitor mixtures under constant driving pressure when diborane was absent from the reactor, and after it was admitted. The difference in flow rates for gas i without and with diborane present, R," - R,, is the rate of disappearance of i due to reaction in the reactor volume. The ratio of competitor flow rates during reaction, Ri/Rj,equals the ratio of their steady state pressures P , / P j in the reactor in the absence of pump fractionation and assuming complete mixing. Total steady-state reactor pressures during reaction varied from -0.009 to -0.35 mm for DhIP-RIP runs, and from -0.03 to -0.28 mm for DMP-TMA runs. Relative rate measurements were made a t 25 2" for 1" for the DMP-TMA both systems, and also a t 9 reaction with cliborane.
RiO= Rt
+ f[Rate of reaction per unit volume]dV
(1)
i.e., the inflow rate must equal the outflow rate plus the removal due t o chemical reaction in the reactor volume, in a steady state. 2. For both sets of competitors, a rough proportionality between relative rate and P a / P jis found, which appears unaffected by a change in the total reactor pressure by a factor of 5 or more. If the runs of common relative rate are grouped together (regardless of total reactor pressure) and plotted against the P,/P5 values averaged within each group, Figures 1 and 2 are obtained. 3. MP and D M P reactions with diborane are of comparable speed, both very much faster than the PH3BzHBreaction, and the infrared spectra of the adducts are almost identical. I
U
Figure 1. Relative reaction rate for methylphosphine (1) and dimethylphosphine (2) competing for diborane, us. relative steady-state concentration ratio a t 25".
Summary of Observations The raw data are somewhat scattered due to the experimental error with which low-pressure flow rate measurements are afflicted, and due to slight temperature variations over the relatively large vacuum system. Significant trends are nevertheless apparent. 1. No significant correlation is found between diborane inflow rate or various functions of this rate (Le., R B ~ H etc.) ~ " ~ and the relative reaction rates for either competitor pair. The diborane flow rate was varied by a factor of 10 for the MP-DMP reaction, and by a factor of 3 for the DMP-TMA reaction. The relative rate rll is computed as (R," - Rt)/(Rj"- R5);a mass balance for each gas gives
I
0
I
&>
I
2.0
Figure 2. Relative reacttion rate for trimethylamine (3) and dimethylphosphine (2) competing for diborane, us. relative steady-state concentration ratio a t 25". (3) Experimental details may be found in the Ph.D. dissertation of W. H. Smith, Department of Chemistry, Syracuse University, 1965. (4) J. Daen and R. A. Marcus, J . Chem. Phys., 26, 162 (1967).
Volume 72, Number 9 September 1968
NOTES
3342
4. Relative rate measurements of TMA-DMP with B2Hsa t 9" under the same flow and pressure conditions as a set of the 25" runs indicate no substantial change in relative rates within the estimated experimental uncertainties.
Contact Angles and Transition
Kinetic Analysis
Lever Brothers Company, Research and Development Division, Edgewater, N e w Jersey 070.90 (Received A p r i l 3,1068)
Because of the nature of the data, relatively little insight can be gained into the mechanisms of these reactions. If one assumes that similar mechanisms hold for M P and DMP-a reasonable assumption in view of the similarity in product structure, reaction rate, and lack of systematic effect of the diborane flow rate on the relative rates-then there are several mechanisms which yield rate equations in accord with the observed behavior, i.e., with
All of these involve the formation of a 1: 1 adduct intermediate, or the formation of a product molecule and BH,, in the first step, but cannot be distinguished by these measurements. If one accepts the mechanism of Bauer, et u Z . , ~ for the TMA-B2H6 reaction in which diborane appears to the first order, then the implication is that the DMP-B~HB reaction is also first order in diborane, since experimentally
Our temperature studies indicate that the relative rates of the TMA-B2H6 and DN!P-B2Hs reactions do not change substantially, implying similar (i.e., low) activation energies. These factors would militate against a mechanism incorporating the diborane-borine equilibrium as an important step. Relative "phenomenological" rate constants can, however, be evaluated without reference to specific mechanisms. Since
Regions in Soap Films
by H. M. Princen
Recently, Mysels, et ul.,' described a technique for measuring the "contact angle" between a thin liquid film and its adjoining Plateau border. The possible existence of such discontinuities had been predicted p r e v i o u ~ l y ~as- ~a necessary consequence of interaction forces in these thin liquid structures. A reported anomaly in the shape of a drop at a liquid-liquid interface6could be explained on this basis.213 Mysels' technique consists of moving an air bubble at a horizontal air-water interface into such a position that the bulk surface is flat up to the circle of contact with the bubble. From the dimensions of the spherical film above the level of the bulk surface, the contact angle can be readily calculated, It seemed to us that the contact angle could be measured more conveniently by utilizing the refractive properties of the Plateau border, which acts to a good first approximation like a prism with finite top angle. Figure 1 illustrates the path of a horizontal light beam t,hrough the region between a vertical flat soap film and the Plateau border which connects the film to the bulk solution. The upper part of the beam passes through the film undeflected, while the lower part is deflected downward. A screen, placed behind the film a t a distance R, shows a bright spot between A and B, a dark region of length D, and a line of light further downward. The angle of deflection is given by tan a = D/R
From Snell's law, applied to the top of the liquid prism, it follows that a is related to the contact angle 0 by a = sin-'
one can write
u n sin 28
Acknowledgment. The authors are grateful to the Research Corporation for supporting a portion of this work. (5) S. H. Bauer, J. V. Martinez, D. Price, and W. D. Jones, "BoronNitrogen Chemistry," Advances in Chemistry Series, No. 42, American Chemical Society, Washington, D. C., 1964,pp 35-52. T h e Journal of Physical Chemistry
- sin-'
}I)%'(
-e
(2)
where n is the refractive index of the bulk solution. For small angles (e), eq 2 reduces to a
V ' Ve find kMp/kD>ip = 0.22 f 0.11, ~ T M A / ~ D M = P 1.60 f 0.44,independent of total steady-state reactor pressure, at 25". Thus ~ T M A : ~ D M P : ~ :8:5:1. MP:
(1)
2(n - i)e
(3) I n real systems the liquid surfaces are bounded by monolayers of stabilizing surfactant molecules. I n N
(1) K.J. Mysels, H. F. Huisman, and R. I. Razouk, J . Phys. Chem., 70, 1339 (1966). (2) H.M.Princen and S. G. Mason, J . Colloid Sci., 20, 156 (1965). (3) H. M. Princen, Ph.D. The& University of Utrecht, Utreoht, 1965. (4) B. V. Deryagin, G. A. Martynov, and Y. V. Gutop, Colloid J . (USSR), 27, 357 (1965). (5) G. D.M. MacKay, Ph.33. Thesis, McGill University, Montreal, 1962.
