THE KINETICS OF THE OXIDATION OF AMMONIA BY NITROUS

OXIDE

By T. N. BELL* and JOY W. HEDGER?

[Manuscript received July 22, 19631

Summary

Ammonia is oxidized by nitrous oxide smoothly and homogeneously at temperatures between 658 and 730' and total pressures up to 250 mm. The products of reaction, nitrogen, water, and hydrazine are accounted for by a free-radical mechanism initiated by oxygen atoms which result from the thermal decomposition of nitrous oxide. Ammonia labelled with the 15N-isotope was used to distinguish between the nitrogen formed from the nitrous oxide and that from the ammonia.

The kinetics follow an empirical rate equation,

Rate = k'[NzO]1'56+lC"[?\TzO]0.G1[NH3].

This is of a form which shows the importance of the ammonia molacule participating in the activation of nitrous oxide through bimolecular collision. Assigning a colli- sional efficiency of unity for like NzO-N2O collision,s, the efficiency of ammonia in the process

N H ~ + N Z O -+ N H ~ + N ~ O * is determined as 0.85 .

Investigations of gas phase oxidation reactions involving nitrous oxide have shown two, clearly defined, type reactions. Firstly, those where initiation is through the thermal decomposition of nitrous oxide to yield oxygen atoms which are then able to attack the substrate, and secondly, those where initiation is through decom- position of the substrate into free radicals which are then able to abstract oxygen from nitrous oxide. The first type of reaction occurs at temperatures above 600" where nitrous oxide decomposes, while the second occurs a t a temperature lower than this, where nitrous oxide is stable. Obviously a mixture of both types may occur when both nitrous oxide and substrate decompose a t measurable rates at similar temperatures.

Examples of type 1 reactions include the oxidation of hydrogen,l carbon monoxide,2 sulphur dioxide,3 and methane,4 while ethane5 and propane6 are of type 2.

* Department of Physical and Inorganic Chemistry, University of Adelaide; present address: Division of Pure Chemistry, National Research Council, Ottawa, Canada.

f Department of Physical and Inorganic Chemistry, University of Adelaide. 1 Melville, H. W., Proc. Roy. Sac. ,4, 1933, 142, 524; 1934, 146, 737. 2 Bawn, C. E. H., Trans. Faraday Sac., 1935, 31, 461. 3 Bell, T. N., Robinson, P. L., and Trenwith, A. B., J. Chem. Sac., 1955, 1440. 4 Robinson, P. L., and Smith, E. J., J. Chem. Sac., 1952, 3895. 5 Kenwright, R., Robinson, P. L., and Trenwith, A. B., J. Chem.Soc., 1958, 660; Kenwright,

R., and Trenwith, A. B., J . Chem. Sac., 1959, 2079. 6 Smith, E. J., J. Chem. Sac., 1953, 1271.

Aust. J . Chem., 1964, 17, 202-11

OXIDATION OF AMMONIA 203

I n the present case the oxidation of ammonia occurs a t temperatures in excess of 600°, and the evidence points to a type 1 mechanism.

A previous study7 of the oxidation of ammonia by nitrous oxide reports no observed dependence of the rate of reaction on the ammonia concentration. This is unlikely for a type 1 mechanism, for i t has been shown8 that foreign molecules may be extremely efficient in promoting the thermal decomposition of nitrous oxide through activating this molecule in bimolecular collision. We have shown that the rate of oxidation is markedly dependent on the concentration of ammonia as well as on the nitrous oxide. Furthermore the ammonia molecules are nearly as efficient as nitrous oxide molecules in transferring energy to nitrous oxide on bimolecular collision.

(i) iiateria1s.-Nitrous oxide, anaesthesia grade, was fractionated through traps cooled to - 72", -90°, - 12S0, and liquid air temperature; the fraction collecting at - 128' was retained.

Ammonia, commercial cylinder grade, was fractionated as for nitrous oxide; the fraction collecting at - 128' was retained. 15N ammonia was obtained in a breakseal ampoule from the Isomet Corp., New Jersey, the isotopic purity was 96%. I t was used directly as supplied.

(ii) Apparatus and Procedure.-A conventional vacuum apparatus was used. The cylindrical reaction vessel was of silica, diameter 36 mm, capacity 163 ml, and fitted with a thermocouple well. The vessel was embedded in a tubular furnace and connected to the rest of the system through capillary tubing, and a ground cone and socket joint. The furnace was maintained to 1 0 . 5 " of the required temperature with a Sunvic type RT2 temperature controller. Temperature measurement was through a chromel-alumel thermocouple in conjunction with a Doran potentiometer.

The reactlon vessel was connected to the nitrous oxide and ammonia storage bulbs through a two-way papillary tap, and also to an analysis train consisting of six traps and a gas burette. Each of the traps could be connected to the gas burette independently. The gas burette was capable of measuring volumes from 0 . 2 to 35 ml wlth an accuracy of 4 0 . 5 % . A constant volume manometer9 was used in conjunction u ~ t h two cathetometers to measure the pressure.

Mass spectrometric analyses were carried out on a A.E.R.E. type S1690 spectrograph, described by Cooke-Yarborough and Russell.lo Gas chromatography was done using a Perkin- Elmer model l54C chromatograph.

The nitrous oxide and ammonia were admitted to the reaction vessel separately, the time interval between successive additions being approximately 4 sec. The static method was used for all experiments involving rate measurements.

(a) Products of Reaction.-The products from a number of runs a t 690°, with initial pressures of ammonia 90 mm and nitrous oxide 90 mm, were divided into condensible and noncondensible fractions.

The noncondensible fraction was examined for hydrogen by attempted com- bustion on copper oxide. No water was formed hence hydrogen was proved absent. Mass spectrometric examination showed the absence of a peak corresponding to mass

7 Volders, A., and Van Tiggelen, A., Bull. Soc. Chim. Belg., 1955, 64, 736. 8 Bell, T. N., Robinson, P. L., and Trenwith, A. B., J. Chem. Soc., 1957, 1474. 9 Bell, T. hT., J . Chem. Soc., 1961, 4973.

lo Cooke-Yarborough, E. H., and Russell, M. C. B., J. Sci. Iwtrum., 1953, 30, 474.

204 T. N. BELL AND JOY W. HEDGER

32; thus oxygen was not present. The noncondensible fraction was therefore assumed to be nitrogen.

The condensible fraction was roughly fractionated through traps cooled to -78", -128", and liquid air temperature. The fraction collecting in liquid air was gas-chromatographed on a 2-m column containing Carbowax 1500 on Teflon support using hydrogen carrier gas. Two peaks were observed, corresponding to nitrous oxide and ammonia. The residue collecting at -78" consisted of water.

In a second series of experiments the condensible fraction was absorbed in 2~ sulphuric acid, and then examined for hydroxylamine and hydrazine. Neither compound could be detected using the method described by Blomll for hydroxylamine, and using the reagent p-dimethylaminobenzaldehyde for hydrazine.12

The products were further examined by conducting a series of flow experiments in which a 1 : 1 mixture of reactants was allowed to flow through a silica tube, dia- meter 3.5 mm, heated to 692" along 20 cm of its length. The pressure was maintained a t 100 mm and a flow rate of 12.5 ml per minute was established. The total products from a run of duration 195 min were passed into 2 . 5 ~ hydrochloric acid. This solution was divided into two parts, and the tests for hydroxylamine and hydrazine applied. That for hydroxylamine proved negative; however, the colour reaction for hydrazine indicated its presence. This observation was confirmed by examining the ultraviolet absorption spectrum of the buffered solution. A distinct absorption maximum a t 458 mp was observed which corresponded to the azine compound12 resulting from the reaction of hydrazine with p-dimethylaminobenzaldehyde.

The products of reaction thus established were nitrogen, water, and hydrazine.

(b) Stoicheiometry.-The stoicheiometry was established from a series of runs at 690" with pressures of ammonia 79 mm and nitrous oxide 91 mm, the time of reaction being varied up to 10 min. The amount of nitrogen resulting from each run was measured in the gas burette, and after separation from the water, the total nitrous oxide and ammonia was similarly measured. The individual quantities of ammonia and nitrous oxide were obtained by gas chromatography of the mixture using a column containing Carbowax 1500 on Teflon with hydrogen as carrier gas.

The nitrogen, as measured, consisted of that from the ammonia and nitrous oxide. This total nitrogen was split into its component parts by repeating the experi- ments with 15NH3, and analysing the nitrogen mass spectrometrically. In this case, in order to obtain sufficient nitrogen for analysis the pressures of ammonia and nitrous oxide were increased to 169 and 341 mm respectively. The mass spectro- metric analyses showed the presence of products with masses 28, 29, and 30, and the absence of masses 31 and 32. The average ratio of mass 28, 29, and 30 for a series of 5 runs was 28(4.14), 29(0.16), 30(1 .00). The small amount of W2 is accounted for by the purity of the 15NH3 (96%).

I t is apparent that the nitrogen atom in ammonia forms molecular nitrogen distinct from that formed from the nitrous oxide.

11 Novak, R., and Wilson, P. W., J. Bact., 1948, 55, 517. 1% Watt, G. W., and Chrisp, J. D., Analyt. Chem., 1952, 24, 2006.

OXIDATION OF AMMONIA 205

The composition-time curve resulting from these experiments is shown in Figure 1. The amount of nitrous oxide lost is in good agreement with the amount of 2*N2 found; thus in the oxidation the overall reactions involving nitrous oxide may be written,

N20 -+ Nz+[O] used for oxidation (a)

If the overall oxidation was written,

3Nz0 f 2NH3 -t 4Nz+3H30

the ratio 2W2 : z9+30Nz would be 3 : 1. Experimentally this has been found to be 3.56 : 1. Furthermore the ratio N20 : NH3 consumed is observed to be 1 .2 : I , whereas (b) would yield a ratio 1 a 5 : 1. Also the loss of ammonia is greater than the amount of 29+30N2 which would result if (b) represented the overall reaction.

Fig. 1.-Composition-time curve. Temperature 690°, ammonia pressure 79 mm, nitrous oxide pressure 91 mm.

However, inclusion of reaction (c) into the stoicheiometry will cause the experimental ratios to be approached.

(c) Rate of Reaction.-The rate of change of pressure with time was taken as a measure of the rate of reaction, the initial rate being obtained from the pressure-time curves.

The effect of the reactant concentrations on the rate was determined a t tem- peratures between 658 and 730". The results where the nitrous oxide pressure was

206 T. N. BELL AND JOY W. HEDGER

varied with respect to a fixed ammonia pressure are shown in Figure 2 , while the dependence of rate on the ammonia pressure with respect to a fixed nitrous oxide pressure is shown in Figure 3,

Fig. 2 Fig. 3

Fig. 2.-Variation of initial rate with pressure of nitrous oxide. Ammonia pressure 57 rnnl

Fig. 3.-Variation of initial rate with pressure of ammonia.

R'itrous oxide pressure 57 mm; 0 n~trous oxide pressure 84 tnm.

It is apparent from Figures 2 and 3 that the experimental rate equation is of the form,

Initial rate = k'(NeO]x+k''[N20]~[NH~]z

The order of reaction with respect to ammonia, z , is, from Figure 3, equal to unity for pressures of ammonia up to twice that of nitrous oxide. Where the pressure of ammonia exceeds twice the nitrous oxide pressure, the order with respect to ammonia falls to less than unity.

The slopes and intercepts of the curves in Figure 3 correspond to kl[NzO]x and ~"[N~OIY respectively. The orders x and y were determined from the ratios of the intercepts and of the slopes for two different fixed pressures of nitrous oxide at a given temperature. The reasonable assumption is made that the orders x and y are independent of pressure for the range of the pressures chosen, namely, 57 and 84 mm.

Values of x and y determined a t 658 and 684" gave x = 1 +56 and y = 0.61.

The experimental rate equation is thus written

OXIDATION OF AMMONIA 207

( d ) Energies of Activation.-k' and k" were calculated from the intercepts and slopes of the plots in Figure 3, making the usual correction for the change in con- centration of nitrous oxide with varying temperature. The Arrhenius plots, shown in Figure 4, yield values of activation energy E' 53.9 and E" 56.5 kcal. Thus E"-E' = 2 .6 kcal.

I I I I I I 101 103 105 107 109

I/T x 105

Fig. 4.-,4rrhenius plot, logl&' and loglolc", against 1/T.

(e) Eflect of Surface.-A reaction vessel packed with silica tubing, diameter of 8 mm, with a surface area/vol, ratio 10 times that of the unpacked vessel showed no variation of rate of reaction for runs a t 692" with pressures of nitrous oxide 57 mm and ammonia 80 mm. The reaction is therefore homogeneous.

If the mode of initiation of oxidation is through oxygen atoms from the thermal decomposition of nitrous oxide, steps 1-5 must be an essential part of the mechanism." It has been shown that the molecular nitrogen formed from the ammonia is distinct from that from the nitrous oxide; step 9 is therefore proposed. The absence of a peak corresponding to mass 31 in the experiments using 15NH3 is important in showing that nitrogen atoms do not abstract oxygen from nitrous oxide.

The following mechanism is proposed to account for the products and kinetics observed.

T. N. BELL AND JOY W. HEDGER

It is recognized that step 8 may not be simple and could conceivably be the sum of several simpler steps; however, rather than speculate on such a mechanism the simpler scheme is suggested. Chain termination is through 10 and 11 ; both aid in accounting for the ratio 28N2 : 29+30N2 observed in the isotopic experiments. Step 10 is undoubtedly reversible with the equilibrium lying towards the left-hand side. Although HOz is a species of only limited stability, its loss at the wall has been used as a chain- terminating step in a variety of oxidation reactions. After a consideration of the possible alternatives Robinson and Smith4 have proposed step 11 as the main chain- ending process in the analogous oxidation of methane by nitrous oxide and show that i t is kinetically and thermodynamically acceptable. In the present case the absence of a decrease in rate in the packed vessel suggests that complete loss of HO2 to the walls takes place before any decomposition into OH and 0 occurs. The above scheme yields a rate expression compatible with that obtained experimentally provided kll[NzO] is small compared with k7[NH3].

Rate of oxidation = dHzO/dt = k7[0H][NH3]+ks[NHz][N20]

Employing the stationary state conditions for the concentrations of 0 , OH, and NHz,

0 = ks[~20c] lke[~~s l , (12)

Assuming

Hence

dH2Oldt = 3ks[~z6]

OXIDATION OF AMMONIA 209

Stationary state for Nzd:

Experimentally

dpldt = ~'[N~O]~'~~+~*[NZO]~'~~[NH~].

This is of the same form as the theoretical equation (18) and is valid for pressures of ammonia up to twice that of nitrous oxide. Above this pressure the order with respect to ammonia is <l ; this is predicted by (18) in that k4[NH3] will be increasingly important with increasing [NH3] for a fixed pressure of NzO.

If (18) and (19) are identified together, a plot of rate against [NH3] as shown in Figure 2 yields the following identities

Intercept = ~ ' [ N ~ O ] " ~ ~ K k l [ ~ 2 0 ~ 2 1 + WNzOIlk.5'

and

Hence

Slope [NzO]/lntercept fi k"/k' = k2/kl.

The relative efficiencies collision per collision of nitrous oxide and ammonia in activating nitrous oxide through steps 1 and 2 may be derived from kllkz. Here the collisional efficiency is defined as the ratio of N20-N20 collisions to N20-NH3 collisions, necessary to cause the same rate of formation of activated nitrous oxide molecules.

From kinetic theory, the collisional efficiency, E , is given by,

where (5 = collision diameter and M = molecular weight.

The collision diameters necessary for the determination of E were calculated from viscosity data using the relationship

2 0 ,499~1 8RT ' (5 =- -

$2 Nn? ( rrM ) ' where p = gas density a t T'K,

7 = gas viscosity a t T'K, N = number of molecules per C.C. = 2.69 x 1019, R = gas constant = 8.314 x 107 ergslmole deg,

2 10 T. N. BELL AND J O Y W. HEDGER

and p ~ , o , (273 '~) = 1.9775 g/l, p x ~ , , (273 '~) = 0.7714 g/l, 7 ~ ~ 0 , (273'~) = 137 /LP, TNH,, (273 '~) = 91 - 8 p P

Hence aN,O = 4'6A; ON*, = 4-5A.

Thus E = k2/kl x 0.76 = kf'/l%' x 0.76.

Values of ktl/k' and corresponding E values are given in Table 1.

TABLS 1

COLLISIONAL EFFICIENCIES OF AMMOTIA

I I

The values quoted8 for other foreign molecules are not strictly comparable with that above owing to the different values of the collision diameter used. However, when the quoted values are recalculated in terms of collision diameters derived from viscosity data, the relative effects may be compared: C02 0 .46 ; CF4 0 a91 ; SO2 0.47; He 0.35; Ne 0.34; Ar 0.18; Kr 0 . 1 ; Xe 0.18.

These efficiency values are for the rare gases at 710°, for SO2 a t 67g0, for C02 mean of values between 651 and 752', and for CF4 mean of values between 673 and 744".

Ammonia is thus extremely efficient in transferring energy to nitrous oxide on bimolecular collision, being far more efficient than the other foreign molecules tested, except carbon tetrafluoride. In no case, however, does the collisional efficiency of the foreign gas exceed that of nitrous oxide itself.

If equation (22) is correct, E"-E' = E2-El. Experimentally this is equal to 2 . 6 kcal. These values are close to those obtained for the thermal decomposition of nitrous oxide in the presence of foreign gases8 where rate expressions similar to (18) and (19) were obtained. This is further evidence that the rate of oxidation is con- trolled by the thermal decomposition of nitrous oxide, where the first term in (18) and (19) represents the decomposition alone, and the second term the activating effect of the foreign gas.

OXIDATION O F AMMONIA 211

The thermal decomposition of nitrous oxide is a complex process; i t has been suggested that, in fact, the decomposition proceeds through triplet states.l3 Owing to this complexity it is quite likely that the small activation energy difference noted is a function of the different triplet states.

We are grateful to Dr. W. W. Forrest, CSIRO, Adelaide, for the mass spectro- metric determinations, and to the CSIRO for a Senior Post-graduate Studentship to J.W.H.

13 Lindars, F. J., and Hinshelwood, C. N., Proc. Roy. Soc. A, 1955, 231, 102, 178.