REACTION OF FORMIC ACIDAND NITROGEN DIOXIDE
3455
Autocatalysis and Deuterium Isotope Effects in the Reaction of Formic Acid and Nitrogen Dioxide
by Donald Barton' and Peter E. Yankwich Noyes Laboratory of Chemistry, University of Ill~nois,Urbana, Illinois
61801
(Received March 19, 1967)
The rates of oxidation by nitrogen dioxide of the hydrogen isotopic formic acids (FA), HCOOH (hh), HCOOD (hd), DCOOH (dh), and DCOOD (dd) have been measured in the centimeter range of partial pressures, at 191.3' in a Vycor vessel, with and without added nitric oxide. The rate law observed is of the form -d(FA)/dt = kl(FA)(NOz) B(NO)*Ia,where B is an incompletely known function of (FA) and (SOz). The finding that there is a small positive intercept on the rate axis when initial rate is plotted lis. initial pressure of nitrogen dioxide suggests the possibility of a small deviation from the indicated simplicity of the nonautocatalytic term. Pollard and Holbrook reported an autocatalytic term of the form k'(FA) (NO,) (NO), but less than first-order dependence on nitric oxide can be seen in plots of some of their results. In experiments with added nitric oxide, a short induction period was observed. Isotope effects on kl are hh/hd = 1.39, hh/dh = 1.48, and hh/dd = 1.99; on B early in the reaction (all experiments at fixed initial pressures of FA and XOZ) they are hh/hd = 1.62, hh/dh = 1.51, and hh/dd = 2.25. The isotope effects are consistent with involvement of both formic acid hydrogens in rate-determining reactions.
+
Introduction Pollard and Holbrook2 studied the reaction between formic acid (FA) and nitrogen dioxide in packed and unpacked Pyrex vessels. They found simple stoichiometry, the products being nitric oxide, carbon dioxide, and water. It was observed that the rate depended upon the partial pressure of nitric oxide and was affected by the surface of the vessel. Their kinetics results were fitted to a rate law of the form given in eq 1 ; rate-controlling bimolecular association was pro-d(FA)/dt = kl(FA)(NOz)
+ k'(F-4) (NOz) (NO)
(1)
posed for the noncatalyzed reaction, and hydrogen abstraction from FA by K203 for the catalyzed reaction. Because the role of hydrogen abstraction reactions in the kinetics of the oxidation of organic compounds by nitrogen dioxide is but imperfectly understood, we undertook to study such processes in at least partial isolation during the oxidation of FA by nitrogen dioxide through the use of the various hydrogen isotopic formic
acids: HCOOH (hh), HCOOD (hd), DCOOH (dh), and DCOOD (dd). The comparative kinetic isotope effects results are reported here, as well as additional findings on the form of the autocatalytic term in the rate law.
Experimental Section Reagents. The ordinary formic acid (hh) used was Baker and Adamson CP grade; formic acidd, formic4 acid, and formic aciddz were obtained from Merck Sharp and Dohme of Canada, Ltd. Samples of each acid were purified using techniques described in the l i t e r a t ~ r e ; ~analyses ?~ by combustion showed the final samples to be pure to about *0.5 mole yo,based on carbon dioxide, and all such were stored a t -78" until aliquoted for kinetics runs. Nitrogen dioxide was purified as in a similar study4 and nitric oxide by the technique described by Nightingale, et aL5 (1) Visiting assistant professor, 1960-1962; Department of Chemistry, Memorial University of Newfoundland, S t . John's, Newfoundland. (2) F. H.Pollard and K. A. Holbrook, Trans. Faraday SOC.,5 3 , 468 (1957). (3) A. S. Coolidge, J . A m . Chem. SOC.,50, 2166 (1928).
Volume 71, Number 1 1
October 1967
DONALD BARTON AND PETERE. YANKWICH
3456
Apparatus and Procedure. The apparatus and kinetic and analytical procedures were essentially as described by Bartone4 The cylindrical reaction vessel was fabricated from Vycor glass; its volume was 203 cm3 and the surface to volume ratio was 1.2 cm-'. The vessel was connected to a quartz spiral gauge used as a null indicator; the pressure was measured to h0.02 cm on an associated mercury manometer. The reactor was enclosed in a heated steel jacket and its temperature regulated at 191.3' by means of a thermistor and electronic relay; temperature uniformity and control were both =kO.l".
Table I : Rate Results, Constant-Rate Region
Results Generally, our experiments confirmed the simple stoichiometry observed by Pollard and Holbrook (though no loss of carbon was found in this study), as well as the principal feature of the plots of pressure vs. time-a section which, within the precision of the pressure measurement (hO.02 cm), appears to be linear, i e . , a region of constant rate. The kinetics runs carried out after a few of exploratory character were of three types, AP referring to the portion of the constant rate region employed in the evaluation of the rate constants, and the pressures being initial values: type A, P F =~ 3 cm, PKO?= various, 0 I AP 5 1 cm; type B, P F A = 6 CM, P N O ?= 3 cm, 0 5 AP 6 0.5 cm; type C, PFA = 3 cm, PNOI = 4 cm, PNO = various, 0.1 AP S 0.8 crn. The kinetics results are displayed in Table I; the values of AP/At listed were obtained from least-squares treatment of all data pairs in the appropriate pressure range indicated above. The data for runs of types A and B are in excellent agreement with the observations of Pollard and Holbrook as to both the magnitude and form of the nonautocatalytic term in the rate law, Icl(FA)(NO2). It was observed in experiments of type B, however, that the initially constant rate increased slowly with time and passed through a maximum near 50% reaction.6 Also, except for dd, there is a small positive intercept on the rate axis when initial rate is plotted us. initial pressure of nitrogen dioxide. The latter finding suggests a small deviation from simplicity in the nonautocatalytic term. The isotopic (kl)i,, equal to (AP/At)/ (PFAPNO.)evaluated with initial pressures and the rate in the constant-rate region, are shown in Table 11, as are certain of their ratios. Preliminary experiments with added nitric oxide indicated that the kinetic order of the autocatalytic term with respect to nitric oxide was definitely less than unity, in contrast to the observations of Pollard and Holbrook.' Further, the period of constant rate is short in experiments \vith added nitric oxide and we obThe Journal of Physical Chemistry
Run no.
Istp FA
9 13 19 20 24 25 26 37 38 39 40 41 42 43 10 12 21 49 50 51 58 61 62 63 11 15 16 17 18 52 53 54 55 56 57
hh hh hh hh hh hh hh hh hh hh hh hh hh hh hd hd hd hd hd hd hd hd hd hd dh dh dh dh dh dh dh dh dh dh dh dd dd dd dd dd dd dd dd dd dd dd
8 14 22 23 27 30 44 45 46 47 48
~
---Init
Run
presa., cm---
FA
NO2
NO
3.02 3.04 3.04 3.04 6.16 6.19 6.10 3.02 3.01 3.04 3.05 3.04 3.05 3.03 3.02 3.03 3.07 3.04 3.03 3.03 6.07 3.03 3.02 3.03 3.01 3.00 3.03 3.01 3.06 3.04 3.02 3.02 3.03 3.02 6.11 3.04 3.02 3.03 6.10 6.12 3.03 3.04 3.06 3.03 3.03 3.03
4.05 11.52 7.21 1.89 3.00 3.03 3.03 4.09 4.08 4.03 4.04 4.07 4.02 3.99 4.07 11.05 7.06 4.04 4.04 4.08 3.03 4.05 4.07 10.63 4.04 13.44 3.97 7.26 1.93 4.08 4.05 4.02 4.04 4.08 3.07 4.04 11.34 1.99 3.07 3.07 4.07 4.06 4.08 4.03 4.07 4.02
0 0 0 0 0 0 0 10.99 5.48 9.03 2.95 0.95 7.00 0 0 0 0 11.81 4.98 1.86 0 12.67 0 0 0 0 0 0 0 9.72 0 5.00 8.67 1.86 0 0 0 0 0 0 0 6.40 9.10 1.00 10.49 3.82
type
A A A A B B B C C C C C C A A A A C C C B C A A A
APlAt, cm/min x 10%
3.75 10.0% 6.12 1.98 5.51 5.72 5.72 17.81 13.28 17.68 10.47 6.50 15.00 3.60 2.85 8.01 4.69 12.92 8.20 5.68 3.71 12.70 2.60 6.77 2.57 -"
A4
P,
A A A C A C
2.50
c C B A A A B B A C C C
C C
I
.an
4 05 1.45 11.94 2.44 8.47 11.39 6.03 3.71 1.92 5.34 0.97 2.77 2.76 1.79 6.48 8.06 2.95 8.27 5.30
~~~
(4) D. Barton, J. Phys. Chem., 65, 1831 (1961). (6) R. E. Nightingale, A. R. Downie, D. L. Rotenberg, B. L. Crawford, Jr., and R. A. Ogg, Jr., ibid., 58, 1047 (1954).
(6) A suggestion of similar behavior may be seen in the plot for Pollard and Holbrook's run no. 17; see Figure in ref 2. (7) Cpnceivably, this discrepancy could arise in the fact that our expenments were carried out at 196O, while the main part of the work of Pollard and Holbrook was done a t 220°, or it might be due to the
REACTION OF FORMIC ACIDAND NITROGEN DIOXIDE
Table 11: Isotope Effects on kl and B" 108(kl)ij , ij
om -1 min - 1
hh hd dh dd
3.03&0.13 2.19i0.11 2.05320.14 1.521:0.06
(kt)hh/(kt)ij
lOaBij , b cm'/s min-1 Bhh/Bij
... 1.39i0.09 1.48=tOO.12 1.99rtO.12
3.15 1.95 2.09 1.40
... 1.62 1.51 2.25
' The Bij are for type C experiments at AP
See eq 3. 0.1 cm.
=
served consistently in such runs a short induction period which has not been reported previously. (This induction period was too short to be characterized well with our manual pressure measuring technique.) Therefore, a series of runs was studied in which the initial formic acid and nitrogen dioxide pressures were fixed a t convenient values, while the initial pressure of nitric oxide was varied. It is reasonable to assume that the simple stoichiometry obtains at all t; then, all instantaneous partial pressures are related simply to the initial pressures and AP. I n the runs with added nitric oxide, and at a particular value for AP, the pressures of formic acid and nitrogen dioxide are identical for each member of the series, but that of nitric oxide is equal to the initial value plus AP. Rat,eswere measured at selected values of AP and were assumed to conform t o the general rate law -d(FA)/dt
=
kl(FA)(KOz)
+X
B = Ic"(FA) (NO,) (4) which decreases throughout the course of the reaction. However, B is a more complicated function of the composition of the reaction mixture than is suggested by eq 4. The plots of the several components of the rate in run 38 which are shown in Figure 2 are typical. Either the order of X with respect to nitric oxide is different from z//3 a t low pressures of nitric oxide, or B is not just a function of (FA) and (NOz), or both;* our results are insufficient for determination of the analytic form of B. Hydrogen Isotope Efects. The isotopic ratios of (k1)ij and Bij in Table I1 span the range 1.39-2.25; such isotope effects are small in comparison with those expected to be normal and primary, but somewhat larger than expected for a normal secondary effectas
.I5
l
l
i
~
i
l
l
~
l
(3)
Values of the isotopic Bij are listed in Table 11, along with certain of their ratios. The correspondence of ea 3 for X with the exDerimenta1 resultsbver a wide in range Of nitric Oxide pressures is shown where the curves are calculated for Af' = 0.1 Cm; the fit is equally good for other AP values.
'!
Discussion d4utocatalys&, Were the difference between our results and those of Pollard and Holbrook confined to
the order with respect to nitric oxide jn the autocatalytic term of the rate law, B would be of the form
l
i
[
-
-
-
(2)
where X is the autocatalytic contribution to the observed rate. Plots of log ( X ) 21s. log (PNo)were made for several values of AP; they are linear, with slope 0.66. The slope equivalent of the scatter in the results is approximately *0.04, except possibly near P = 1 cm, and there is no drift of the mean slope with AP. Thus, to a good approximation
x = B(NO)*'/"
3457
-
0
4
8
12
Pw(cm) Figure 1. Comparison of X o s l c d and Xobsd as functions of PNOfor each isotopic FA; calculations are for AP = 0.1 cm. Initial pressures of FA and NO, are 3.03 and 4.05 cm, respectively; the following symbols represent the isotopic formic acids: 0, HCOOH; @,HCOOD; o, DCOOH; and ., DCOOD.
20% difference between reaction vessel surface to volume ratios in the two studies. However, examination of their Figure 3 (a plot of initial rate vs. initial pressure of nitric oxide) shows that in-every case a curve through the data points bows upward with respect to the line drawn through them, a clear indication of an apparent kinetic order with respect to nitric oxide of less than unity. ( 8 ) Other experiments (part of a different study and not reported here) show that B is not symmetrical in nitrogen dioxide and formic acid pressures; reversal of the initial values of the two causes a pronounced change in the shape of the P-t curve. Also, the plot of log ( X ) vs. log (PNo)suggest that there may be significant deviation from two-thirds order a t low pressures of nitric oxide.
Volunx 71, Number 11 October 1967
~
DONALD BARTON AND PETER E. YANKWICH
3458
.I 2
.08
.04
0 I
0
2
3
ARcm Figure 2, Variation with AP of the quantities R, total rate; A , noncatalytic contribution to observed rate; X , autocatalytic contribution to observed rate; and B, = X/(NO)e/a; run no. 38, type C.
Barton4 has shown how a modest isotope effect can be multiplied in a complicated oxidation superficially similar t o that studied here. It is easy to show that such reasonable mechanistic features as dynamic opposition of chain initiation and termination steps in the oxidation of formic acid by nitrogen dioxide can yield a small apparent isotope effect when all of the isotope
The Journal of Physical Chemistry
effects associated with the individual elementary reactions are of normal magnitude. This reaction has a strong heterogeneous component; however, if strong adsorption of formic acid on the vessel surface resulted in preferential attack at one hydrogen (attack on C-H is favored energetically, and the adsorption would likely be strongest through 0-Ha . 0 hydrogen bonding), ( k l ) h d and (k1)dh would not be expected t o be so similar. The similarity of the hd and dh isotope effects indicates that changes at both 0-H and C-H are important t o the mechanism of the reaction; however, we have no explanation for the fact that (kl)hd > ( k l ) d h while B h d < B d h . Detail of the NO Catalysis. While our results do not exclude Nz03as the catalytic agent, they do exclude a mechanism as simple as that proposed by Pollard and Holbrook. No doubt, KO radical addition reactions and hydrogen abstraction by nitric oxide should be among additional steps in the mechanism. It is of interest that catalysis by nitric oxide occurs also in the nitrogen dioxide oxidation of acetic acid;*O there may be a relation between the observation of catalysis and the presence in a reagent of a carboxyl or carboxylrelated function. Further study of the reaction between nitrogen dioxide and formic acid should contribute to understanding of the role of nitric oxide in such oxidations.
Acknowledgment. This research was supported by the U. S. Atomic Energy Commission, COO-1142-73. (9) L. Melander, “Isotope Effects on Reaction Rates,” Ronald Press Co., New York, N. Y . , 1960, Chapters 4 and 5 . (10) D. Barton and R. N. Pandey, in preparation.
