298 TABLE I11 THEAPPARENTDISSOCIATION C O N S T ~ KOF T SC a ~ c rair IWDOPHOSPH ~TC\ Calcium cornplev of
Temp.,
Iinidodiphosphatt-1
25
5 59 f 0 5 3i f 5 14 f t 3 3 3 f G 74 It F 21 i 0 01 i. 1 44 f
3i
11onohydrogen imidodiphoqphatc I,liiniidotripho.phatf,
11onohrdrogen diimidotriphosphate C H 4NGES I N
log of dissociation constant a t ionic stlengths of-03 10
---Negative 01
OC.
50 25 25 37 50 25
07 02 04 10 12 05 08 10
TABLE IT' THERMODYNAMIC QUANTITIES FOR C 4 L C I I W
Calcium complex of
Negative log of T e m p , , thermodynamic molal "C. dissociation constant
5 20 f 0 05
4 59 It 0 4 14 i 3 8G f .i15 f 5 GG =t 51Li. 4 86 It 4 16 It
G 1 1 i . O OF
09 09 04 10 12 10
10 IO
COMPLEXISG BY IiVIDOPHOSPHATES [ A F a ] , " ?b kcal.
AH", kcal.
[ A S o ] , a >b B.U.
25 607f01 - 5 6 f O 1 -57+00 0 f3 37 5 8 9 i I - 6 l f 1 50 5 7 6 f 1 - 6 1 5 1 25 3 4 f 2 -222t 2 AIonohydrogen imidodiphosphate Iliimidoi riphosphate 25 7 1 o i 2 - i 4 f 2 -9 l f 2 -6 i7 37 6 9 0 5 2 - 7 4 f 2 50 660& 2 -74i 2 Monohyirogen diimidotriphosphate 25 4 6 f L -39f 2 * Taking the standard state as a hypothetical unity mole fraction. b The influence of temperature was taken into acroiint in applying the corrections to obtain the unit mole fraction standard state.
Imidodiphosphate
phate yields values of 3.1 and 3.5 .&., respectively, so that bonding by this criterion is ionic rather than covalent. The second possibility is to assume that waters of hydration were released, together with the forma-
tion of a more ordered structure, with both entropy change contributions cancelling one another. Acknowledgment.-The authors thank Mr. TT' W. Morgenthnler for making qome of the measurements.
OXIDATION OF ,4MMONIL4 IS FLAMES BY C. P. FENIMORE ASD G. W.JONES General Electric Research Laboratory, Schenectady, A'. I'. Recezved J u l y $8 1960
The rate of the reaction of nitric oxide with ammonia has been measured in lox pressure flat flames a t 1'700 to 1900'K. by probe sampling, and simultaneous estimates of [HI obtained. The data support the vim- that the equilibrium K[?;H?]-
+
- I,
[H2]= [ SH3][HI is satisfied for small ratios of [NO]/[H,], and that nitric oxide decays vau reaction I, KO YH2 . , with k l / K = 6 X 1010 1. mole-' sec.-l. Reaction 1 is important in other, suprrficially different, deeventually K2 compositions of ammonia also. Much nitric oxide is generated in the reaction zone of H2-NH1-02 flames, probablv by attack of 0 atoms on ammonia, and nitrogen is formed chiefly by reaction 1. Furthermore, residual ammonia siirvives into the post-flame gas of rich enough ammonia flames and decomposes there, in the absence of appreciable oxygen or nitric oxide, much more slowlp than in the reaction zone; and this too map he due to a slow formation of transient nitric oxide follomd b j rapid 1.
+-
Introduction Adams, Parker and Wolfhard' showed that nitric oxide decays faster in ammonia than in hydrogen flames, and since nitric oxide rapidly scavenges NH?radicals according to kl
NO
- NH2 +
+eventually N2 +
(1)
even a t room temperature123 they suggested that the same process occurs in flames. I n this paper we show that the decomposition of ammonia in (1) G. SGC, 14, (2) C. (3) A (19.59)
K A d a m ? , W. G. Parker a n d H. G. Wolfhard, Dmc. Paraday 97 (1953) H . Barnford, Ibzd., 38, 568 (1939). Sereuicz a n d W. Albert Noyes, Jr., J . Phys. Chem., 63, 843
many flames involves reaction 1. If nitric oxide is not supplied as a reactant, it is generated by oxidation of part of the ammonia and then consumed via reaction 1. Adams and co-workers considered that R reNO + N2 OH, might be imaction, S H portant as well. But under the conditions we used, it was unnecessary to assume an important role for S H radicals. Experimental
+
+
A water-cooled, porous, flat flame burner,d mounted in a bell jar with movable quartz probe and quartz coated thermocouple, was fed with reactants metered through (4) W. E. Kaskan, "6th Symposium on Combustion,'' Reinhold Publ. Corp., New York, N. Y.,1957, p. 134.
OXIDATIOX OF AMMOSIB I N FLAMES
Feh., 1961
critical flow orifices. Commercial nitric oxide from cylinders contains a trace of nitrogen dioxide and, when mixed with ammonia, forms a solid which tends to plug the porous burner, but this difficulty was overcome by passing the nitric oxide at pressures of several atmospheres through an aluminum coil, 1 m. long by 3 cm. i.d., immersed in a Dry Ice trap. Gas samples, taken through the fine quartz probe into a samule bottle maintained at 1 mm. -pressure or less, were inalyzed on a mass spectrometer.
Estimating [HIin the Reaction Zone.-Reaction 1 was studied by adding small amounts of ammonia and nitric oxide to lorn pressure, fuel-rich NZO-H, flames. In these flames, nitrous oxide deconiposes mostly by reaction with H atoms a t a known rate6 -d[K,O]/dt= 4 X 10l1e-I63'RT [HI [rS,O] mole 1.-l sec.-l and therefore [HI can be estimated in the reaction zone. Figure 1 s h o w some profiles through a typical flame, without added ammonia or nitric oxide, and illustrates the measurement of [HI. tTsingthe diffusion coefficient plotted, me recast the nitrous oxide profile as the curve for -d[N20J/dtat the bottom of the figure15and derive [HI values which approximately double through the reaction zone. The increase may be greater than we have found in other low pressure flames; [HI determined in the same way in mixed HPair-SzO flames appeared more ~ o n s t a n t . ~ The constant for the nitrous oxide reaction is based on measurements of [HI via the exchange reaction H D20 = H D OD; and although such measurements parallel results obtained in other ways, for example from Kaskan's optical measurement of [OH] in the same laboratory,6 they are smaller in absolute value by about a factor of' two. ,1 systematic error, if one exists, could not affect relative [HI, however, nor a derived rate constant relative to the constants for the H N2O or H 0 2 reactions as given in this paper. Checkfi of [HI by exchange reactions with D2 or D20bin about half the runs showed that the value of the rate constant given above for H S20+ X2 OH remained unchanged. Since the reaction of added nitric oxide was studied in flames like that plotted in Fig. 1, we should state that a little nitric oxide was formed in this flame. The generated nitric oxide, however, was less than 1/10 of the nitric oxide decomposed in the studies to be described in the next section. Rate of NH2 NO -t . . .-+ N2 . .-Figure 2 shows some profiles through the same flame as Fig. 1 except that the pressure was raised to 12 em. and additions were made to the reactants. For the dotted curves, 0.15 mole of nitric oxide was added per mole of nitrous oxide fed; for the solid curves, the same nitric oxide and also 0.12 mole of ammonia was added. The nitric oxide survived when ammonia was absent, but largely reacted when ammonia was present. At 0.6 cm. from the burner surface, 74y0 of the added nitric oxide and almost all the added ammonia had disappeared; so the reaction proportions were approximately equimolar.
+
299 .zoo0
0 2,-
$ "
1 Y
r
-1500
= *
OIL
5
* Y
I &
_____
1
03 06 DISTANCE FROM BURNER SURFACE.
+
+
Fig. 1.-Profiles through a flame of ?jzO 1.6 H, 2.32 A burnt at 6 cm. P with a mass flow of 3.56 X 10-3 g. cm.-2 sec.-l.
+
+
+
0
+
+
+
+ .
( 5 ) C. P. Fenimore and G . W. Jones, J. P h y s . Chem., 63, 1154 (1959). (0) W. E. Kaskan, Combustton &: Flame, 2 , 286, 229 (1958).
