3266
JOSEPH
H. DUSENBURY AND KICHARDE. POWELL
photometer (model U U ) it1 the region 216 mp to 360 inp; E = ( l / c ) log Io/I, where c is the concentration in moles per liter. Siiicc most of the compounds were rather unstable liquids, a t least three distillation cuts of similar boiling point arid refractive index were uscd and the triplicate spectral determinations (in 95% ethanol) wrre carried out imniediatcly after distillation. Preparation of Compounds.-In spite of the fact that most of the substances are known, their purity was checked by elementary analysis in addition to b.p. and refractive index. Compounds 1, 3 , 16 and 22 were prepared according to Sield,ls X o . 5 according to Pauly and Lieck,lg Nos. 17 and 19 by decomposition of the corresponding Mannich bases,20 while S o s . 20 and 21 were gifts of Dr. Jay S. Buckleyof the University of Minncsota.ZL The synthesis of diacetylbutadiene ( N o . 8)22will be published a t a later date.23 The following compounds are new: 3,6-Dibromo-3,5-octadien-2,7-dione (No. 7).-This compound was obtained in 56% yield from the corresponding d i l a c t ~ t i eby ~ ~ following the gcneral method described by Nicld.18 Two recrystallizations from 95% ethanol afforded the desired compound in rosettes of yellow needles, m.p. 102- 103 '. .~lital. Calcd. for CIIIH802Br~: C, 32.46; H, 2.72. Found: C , 32.60; H, 2.79. .. ~
(183 C 11.Nield, THI?J O U R N A L , 67, 1145 (1943). il!li H. Pauly and H. Lieck, Ber., 33, 500 (1900). (20) K. Mannich and C . Heilner, ibid.. 66, 356 (1!)22j; K. Matinich aiid I) Lammering, i b i d . , 55, 3510 (1922). [ ? II K. T. Arnold, J. S. Ruckley and J Kichter, 'THIS J O U R N A L , 69,
Vol. 7 3
3-Brom0-4-(p-methoxyphenyl)-d-buten-2-one (No. 24).a-Keto-p-acetyl- -/- (p-methoxypheny1)- butyrolactone % a i synthesized in 32% yield by exactly the same procedure employcd by Nie1dls for the corresponding -/-phenyl derih citive; 1n.p. 157-158". .{nul. Calcd. for CUHI~O.,: C, 62.90; H, 4.88. Found: C, 63.17; H, 5.29. The bromination arid cleavage of the butyrolactone derivative was carried out according to Nield's method18 a t 5" in 85% methanol solution; yield 45010, b.p. 153-155' (0.01 mm.), n ' 9 ~1.6579. .4naZ. Calcd. for CIlH1102Br: C, 51.79; H, 4.35. Found: C, 51.56; H, 4.62. a-Bromovinyl-p-methoxyphenyl Ketone (No. IB).-The synthesis of a-keto-6-(P-methoxybenzoy1)-butyrolactoneas modeled after that of the corresponding benzoyl derivativeL8; yield 34%, m.p. 153-154'. Anel. Calcd. for C12HloOj: C, 61.54; H, 4.30. Fouiid: C, 61.66; H , 4.66. The bromo derivative was prepared from the above lactone in the usual manner18 in 41% yield, b.p. 130-131' (0.005 mm.),n 1 9 ~1.6028. A n a l . Calcd. for ClnHgO.Br: C, 49.82; H, 3.76. Found: C, 49.71; H, 3.99.
122) P. Karrer. C. H Eugster and S. I'erl, lirirl. C / z i u ~ . Acta, 32, 101:1. 1!131 (1049>, 1;.Soiidheimer a n d ! F C I,, \Veedon, hruti~yc,166,
Acknowledgment.-The authors are greatly indebted to Srta. Paquita Revaque for assistance in the spectral determinations and to Prof. A. L. Wilds of the University of Wisconsin for his cornments.
,482 11950j. c2Sj R. B. Woodward and K. Daniels; Ilarvard University, 1950.
MEXICOCITY,L). F. CAMBRIDGE, MASS.
[CONTRIBUTION FROM
THE
G/.
K. Daniels, 1'h.D. Thesis,
DEPARTMENT OF CHEMISTRY
AND
CHEMICAL ENGINEERING OF
RECEIVED JANUARY 2 , 1951
THE
UNIVERSITY
OF
CALIFORNIA]
Reactions of Nitrous Acid. I. Ammonium Nitrite Decomposition1 B Y JOSEPH
H. DUSENBURY AND
RICHA4RDE. POWELL
The rate law for the reaction of nitrous acid with ammonia is d(iYn)/dt = k, (NHa+)(HNOz), where k, = (kT/k) exp (-7.75/R) exp(-ZO,OOO/RT) sec.-I mole-' liter. The rate-determining step of the reaction is proposed t o be reaction of iiitrosyl ion with molecular ammonia.
The kinetics of the reaction between ammonium ion and nitrite ion SH4'
+ SOn- = S. + 2HzO
was studied by Abel and co-workers in 1931.' They reported the reaction to be kinetically third order: first order with respect to ammonium ion, to nitrite ion and to molecular nitrous acid. Reaction mechanisms involving such an extra molecule of nitrous acid have seemed to us implausible, so we have reinvestigated the kinetics. Experimental The reaction vessel was a three-necked flask equipped with a gas-tight motor ~ t i r r e r ,a~ pressure-equalizing funnel for adding the sodium nitrite solution t o the other reactants, arid a gas buret in which the evolved nitrogen was measured. The reactor was immersed in an oil thermostat whose temperature was controlled to 10.05'. The apparatus was swept with nitrogen before each run. At t h e completion of each run the evolved gas was shaken with permanganate solution to determine the nitric oxide liberated by the in(1) Presented a t t h e 115th National Meeting of t h e American Chemical Society, San Francisco, California. (2) E . Ahel. H. Schmid a n d J . Schafranik, Z . g h y r i k . Chem., Bodenstein Festband, 610 (1931). ( 3 ) \V. G . Dauben, J C Reid arid 1'. E . Y , ~ i i k ! r i c l i ,A tiill C h m . , 19, %28 (1947)
ciderital decomposition of nitrous acid. Under our experimental conditions the amount of such decomposition was negligible. Reaction mixtures were made up from stock solutions of sodium nitrite and ammonium perchlorate, whose concentrations were checked by analysis. The acidity of t h e reaction mixtures was controlled by phosphate buffers made with phosphoric acid and NaHtPOd, and was measured before and after each run with a glass-electrode pH meter. The ionic strength was brought to the desired value with sodium perchlorate. Individual measurements of initial rates were reproducible within about 2%.
Results and Discussion m e series of rate measurements, only the initial concentration of ammonium ion was varied (Table I , Fig. 1). In another series, only the initial concentration of total nitrous acid was varied (Table 11, Fig. 2 ) . These data demonstrate that the reaction is first order with respect to (stoichiometric) ammonia and first order with respect to (stoichiometric) nitrous acid. We find no evidence of the second-order dependence on stoichiometric nitrous acid which Abel, et al., reported. The probable explanation of the discrepancy is that we buffered our reaction mixtures, whereas they 111
omitted doing so.
July, 1951
AMMONIUM NITRITEDECOMPOSITION
TABLE I CONCENTRATION^ EFFECTOF AMMONIA ZHN02, A 4
ZNHa, IM
4
Initial rate, mole liter-' sec. - 1 x 10-
2.86 0.00904 0.395 128 2.87 ,00896 ,197 64 2.82 ,00916 ,098 34.9 2.84 ,00924 ,049 16.6 a Temperature 30.0", ionic strength 1.OO, total concentration of phosphate buffer 0.567 M .
TABLE I1 EFFECTOF TOTALNITROUSACID CONCENTRATION^ Initial rate, mole liter - 1 sec. - 1 ZHNOs, M
#H
ZNHs, M
3267
x
10-
0.0940 0.186 643 2.98 2.92 ,0507 .196 338 ,196 335 2.90 ,0488 2.91 ,0249 .196 156 2.91 ,0100 ,198 65.0 .198 32.6 2.94 ,00490 2.90 ,00243 .198 17.5 a Temperature 30.0', ionic strength 1.OO, total concentration of phosphate buffer 0.567 M .
The dependence of the rate upon $H (Table 111, Fig. 3) permits us to decide which of the NO?- or various nitrous acid species ("02, possibly NO+) may react with which of the ammonia species (NH4+, NH3 or possibly NH2-) in the rate-determining step. For comparison
a,
E +-' w
e
0.01 0.1 1 Stoichiometric concii. of ammonia, ill. Fig. 1.-Dependence of rate on ammonia concentration The straight line has a slope of unity. 10-6
d
si i
-
'k
x
a,
3 10-6
E
c) a,
e
TABLE I11 EFFECT O F HYDROGEN I O N CONCENTRATION' Initial rate, mole liter-' sec. - 1
flH
ZHNOz, 1l.P
ZNHa, M
x
10-7
10-8
0.43b 0.0477 0.197 517 ,196 513 0.96 ,0475 ,193 371 1.81 ,0439 2.95 .0481 .197 305 ,197 140 3.68 ,0485 ,196 51 7 4.11 ,0476 4.33 0487 .I98 30 5 5.00" ,0480 ,197 9 15 6 . lgd 0492 ,198 0 588 a Temperature 30.0', ionic strength 1.OO, total concentration of phosphate buffer 0.567 M . * Acidified with sulfuric acid. Rate was measured at 43.0' and recomputed to 30.0" on basis of 20 kcal. heat of activation. Rate was measured at 60.0' and recomputed to 30.0" on basis of 20 kcal. heat of activation.
with the experimental data, we compute a family of curves giving the expected rate as a function of p H for all possible pairs of reactant species. The ratio of the rate a t any given acidity, (H+), to the rate a t 1 M H + is given by the expression
where 2 is the charge on the activated complex, which is the sum of the charges on the two reactant species. K1 and K2 are the known ionization con,~ stants of HNOZ and NH4+, r e s p e c t i ~ e l y . ~The several theoretical curves are displayed in Fig. 4. (4) A. Klemenc and E. Hayek, Monalsh., 54, 407 (1929). (5) D . H. Everett and W. F. K . Wynne-Jones, Proc. R o y . Sac. (London), 8169, 190 (1938).
0.001 0.01 0.1 Stoichiometric concn. of nitrous acid, M . Fig. 2.-Dependence of rate on nitrous acid concentration. The straight line has a slope of unity.
a .I . YI
4
10-9 0
1
2
3 4 5 6 7 PH. Fig. 3.-Dependence of rate on acidity. Circles, measurements a t 30". Squares, measurements a t higher temperatures recomputed to 30'. Crosses, measurements by Abel et d., computed from data in ref. (2). The curve is the theoretical curve for 2 = +1.
The experimental data (Fig. 3) are in excellent agreement with the theoretical curve for 2 = +1, but in no kind of agreement with any of the other curves. The activated complex for the reaction therefore carries a +1 charge. I t has the formula of the nitrosoammonium ion NHBNO+ or a hydrated form thereof, and could be formed
JOSEPH
subject to general acid catalysis by the phosphoric acid species. At higher ionic strengths the reaction is noticeably slower (Table IV), as Abel, et al., observed also. This phenomenon is surprising, for we would
104
-
103
>
10L
Vol. 73
H.DUSENBURY AND RICHARD E. POWELL
z 74
TABLE IV EFFECTOF IONICSTRENGTHa
I i
-;
10
4
v c
A
1
9H
ZHsPOd, M
sec.
-1
Ionic strength
L-( b+
-
3.0 0.226 3.0 ,226 3.0 ,226 2.0 ,576 2.0 ,567 2 0 ,567 a Temperature 30.0'.
Q,
.-M
0.1
#
2 10-2 0
4
810
3
10-4
0
1
2
:
