D. BEHAR,D. SHAPIRA,AND A. TREININ
180
Values of Eoare plotted below those of rlT in Figure 6. For (py)/(chl) = 19, the extrapolated value rlT = 150 was used. The effect of association of pyrochlorophyll on Eo is readily seen. Instead of increasing coincidentally with TIT,the curve for EOreaches a plateau value of about 8, then increases again as pyrochlorophyll becomes predominantly associated. The plateau value represents a limitation imposed by competition with trapping by aggregated pyrochlorophyll molecules. Increased efficacy a t high pyrochlorophyll densities corresponds to transfer of energy through the shortlived states of aggregated pyrochlorophylls to the even shorter lived pyrochlorophyll-DTNB traps. (In these calculations, back transfer from pyrochlorophyllDTNB to associated pyrochlorophyll was neglected; inclusion of it would appreciably reduce the efficacy of trapping a t high pyrochlorophyll density.) In 1964 Duysens remarked that here twas little evidence for chlorophyll dimer formation in the chloro-
plast, and that one of the functions of the lamellar structures might be to keep the molecules of chlorophyll apart.la The cogency of that remark is well illustrated in our model systems, where it is evident from Figure 6 that "photosynthetic units" 10 times the size of the calculated ones could have been attained if the self-quenching of pyrochlorophyll could have been avoided. Reaction center chlorophyll generally absorbs a t longer wavelengths than antenna chlorophyll. Whether this plays a necessary role in the transference of energy to the trap, or is only an epiphenomena1 consequence of the special environment surrounding reaction center chlorophyll, is unknown. DTNB certainly is not a natural constituent of the chloroplast, but its association in a complex with pyrochlorophyll indicates how an electron accepting natural constituent could exist in a complex with chlorophyll that is stabilized by attachment of both to the lamellar substrate.
Acknowledgments. This work was supported in part by National Science Foundation Grant No. GB-7893. The able technical assistance of T. H. Meyer is appreciated.
The Photolysis of Hydroxylamine in Aqueous Solution by D. Behar, D. Shapira, and A. Treinin* Department of Physical Chemistry, Hebrew University, Jerusalem, Israel
(Received August 18, 1971)
Publication costs borne completely b y The Journal of Physical Chemistry
The photolysis of NHzOH in aqueous solution a t 2139 8 can be closely represented by the overall reaction with some H, also produced. Flash photolysis and pulse radiolysis of hyNZ 3NHzOH NHI droxylamine solutions have proved t h a t NHOH is a n intermediate in the photolysis, and the properties of this radical were investigated. The effect of ethanol on the yields of NHOH and the final products strongly
+ +
hv
+
suggest t h a t the major primary process is N H 2 0 H -.L NHz OH. A minor parallel process leads t o the production of Hz. The photolysis of NHaOH+ was also investigated. At p H -2 the major nitrogen-containing products are NHa and NzO. This change is related to transformation of the hydroxylamino radical to its acidic form (pK = 4.0 f 0.1). Additional information is presented on the reactions of hydroxylamine with NH20H) = 6.6 x 10s IM-1 sec-l, k(e,,NHaOH+) OH, H, and ea"-, The rate constants derived: Ic(e,,= 1.0 X 10'0 1 M - I sec-l.
+
The direct photolysis and Hg-photosensitized decomposition of NHzOH has been studied in the gas phase.l!z According to Betts and Back2 the direct photolysis involves two primary processes yielding NHz 3. OH (40%) and NHzO H (6001,), respectively, whereas the Hg-photosensitized decomposition proceeds primarily by the first process. The radicals NH2, OH, and H were considered subsequently to attack
+
The Journal of Physical Chemistry, Vol. 76,No. 0,1972
+
NHzOH yielding NHa, HzO, and H2, respectively, as final products and the intermediate NHOH, the annihilation processes of NHOH and NHzO being the source of Nz and NzO,respectively.2 Some of these re(1) R. N. Smith and P. A. Leighton, J . Amer. Chem. floc., 66, 172 (1944). (2) J. Betts and R. A. Back, Can. J . Chem., 43, 2678 (1966).
PHOTOLYSIS OF HYDROXYLAMINE IN AQUEOUS SOLUTION actions were postulated for the radiation chemistry of hydroxylamine in aqueous solutionS and some information on NHOH has been recently obtained from pulse radiolysis experiments. Here we present the first results on the photochemistry of hydroxylamine in aqueous solution. Both steady and flash techniques were employed for this purpose. Additional information on the formation and properties of NHOH were also derived from pulse radiolysis of aqueous hydroxylamine. Experimental Section
A . Steady Photolysis. Solutions of hydroxylamine were irradiated a t 23 f 1” with a 25-W Philips zinc lamp No. 93106; the light was filtered by 0.5-cm layer of water. Only the strong 2139-A and weak 2062- and 2025-i lines were absorbed by the solutions. Cutting off the weak lines with a KC1 solution, we could show that about 96% of the photolysis was due to the 2139-A line. The optical path of the irradiation cell was 3 cm, and it was attached to a 25-ml bulb, where the solution was degassed before irradiation. Chemical actinometry was conducted with azide soM NaNa lution.6 An air-containing solution of at p H 7.7 (with a borate buffer) was irradiated under the same conditions and the azide depletion was deter= 0.32 f 0.01. mined, taking Qj-~a The degassing and analysis of products were conducted as described elsewhereS6 The gaseous products were also subjected to mass spectrometric and gas chromatographic analysis, using appropriate blanks. B. Flash Photolysis and Pulse Radiolysis. The flash apparatus and experimental technique were described elsewhereS6 A linear accelerator operating a t 5 MeV and 200 mA was used for pulse radiolysis. The pulse duration was varied in the range 0.2-1.5 psec in order to change its dose. An optical path length of 12 cm was attained by passing the monitoring light three times through the irradiation cell. A modified Fricke dosimeter’ was used to determine radical yields. To reduce stray light effects a 60-W deuterium lamp was the source of monitoring light and for detection we used a “solar blind” R166 photomultiplier, which does not respond to light above 320 nm. With the aid of two filters we measured the stray light coming from within and outside the absorption region of the transient. The first contribution was predominant and as an approximation its intensity was assumed to be reduced by one-half owing to absorption. C. Materials and Solutions. Solutions of hydroxylamine were prepared from (NH20H)2 HzS04 (UCB, analytical grade). I n the photochemical experiments (steady and flash) the pH was adjusted in the range 1.0-6.7 by adding HClO, or NaOH; pH 7.4 was attained by neutralizing the material with an equivalent amount of NaOH and adding a borate buffer. I n the pulse radiolysis experiments, when studying reactions
181
of eaq- produced by 0.2-psec pulses, self-buff ering of the hydroxylamine system (in the pH range 5.4-6.7) and a borate buffer were employed. I n other pulse experiments a phosphate buffer was used for the range above pH 4.0. Ethanol was of spectroscopic grade, water was triply distilled, and all other materials were of analytical grade. Results A . Steady Photochemistry. Figure 1 shows the uv absorption of hydroxylamine at pH 6.9, where ca. 90% is in the basic form. The acidic form absorbs much less in this region. The band of NHzOH in water appears to be blue shifted by -7 kcal relative to the gas phase.2 This energy is comparable to the solvation energy of gaseous NHzOH (AH,,,1 = -11.5 kca18) and therefore similar to other related molecules the assignment of the band to n + u* transition seems p l a ~ s i b l e , ~ with the nonbonding electrons of N being mainly in-
405
IS0
1%
WO
--J
_
,
2005
El0
215
~
0
L,nm
Figure 1. The electronic spectrum of NHzOH in aqueous solution a t pH 6.9. (3) M. Lefort and X. Tarrago, J.Inors. Nucl. Chem., 16, 169 (1961). (4) M. Simic and A. Hayon, J. Amer. Chem. SOC.,in press. (5) I. Burak, D. Shapira, and A. Treinin, J . Phys. Chem., 74, 568
(1970).
(6) D . Behar and G. Czapski, Israel J. Chem., 6, 43 (1968). (7) M. S. Matheson and L. M. Dorfman, “Pulse Radiolysis,” MIT Press, Cambridge, Mass., 1969. (8) R. A. Back and J. Betts, Can. J. Chem., 43, 2157 (1965). (9) D . P. Stevenson, G. M. Coppinger, and J. W. Forbes, J. Amer. Chem. SOC.,83, 4350 (1961).
The Journal of Physical Chemistry, Vol, 76,No. 2, 1972
182
D. BEHAR,D. SHAPIRA, AND A. TREININ
Table I: Yields of Photolysis in Air-Free Solutions of Hydroxylamine at 2139 A -Solution[NHzOHl,
M
2 x 10-3 2 x 10-3 2 x 10-3 2 x 10-3 2 x 10-8 2 x 10-8 2 x 10-3 2 x 10-3 2 x 10-3 2 x 10-8 1.0 1.0 5 x 10-2 5 x 10-8
[Ethanol], M
... ,..
... ...
... ... ... 0.06 0.43 0.43 0.10 .
I
.
...
...
PH
Time, min
7.4 7.4 7.4 7.4 7.4 7.4 7.4 7.4 7.4 7.4 7.4 7.4 4.3 2.0
72 100 107 116 120 150 150 140 140 140 45 45 180 180
--
10-6 My--------
--yield,
2Wzl f
-
[Nrl
IHzl
[NHa]
-A[NH20Hla
[NHal
[NzOl
14.5 19.0 21.3 22.5 22.2 26.0 30.0 8.6 5.4 5.3 18.2 17.7 13.4 2.1
1.6 1.9 2.3 2.6 2.8 2.9 3.1 2.9 3.3 3.2 1.9 1.9 0.6
17.0 18.6 20.5 23.0 23.0 26.2 27.0 22.8 23.0 23.0
47.2 58.0 64.0 67.8 70.3 81.5 85.2 41.3 34.3 33.9
46.0 56.6 63.1 68.0 67.4 78.2 87.0 40.0 33.8 33.6
...
...
... ,
I
.
... ... ...
...
2.4 10.6
-3Ob
In presence of alcohol acetaldehyde is probably formed (from the alcohol radical) and reacts with NH20H to produce the oxime. However, under the conditions employed for the analysis (ref 5), all the oxime was hydrolyzed. 3. In presence of 5 X lo-$ M hydroxylamine the determination of ammonia is not accurate: estimated limit of error, &15%.
volved. The large decrease of absorption on protonation is also in agreement with this assignment. The photolysis of NHzOH in aqueous solution (pH 7.4) leads to the fortnation of nearly equivalent amounts of Nz and NHa, as major products, and a relatively small amount of Hz. No NzO could be detected: all the gas that could be condensed at liquid air temperature could also be condensed in a COz-acetone trap.1° Working with 2 X M NHzOH we measured the yields of products and depletion of NHzOH; the results are summarized in Table I and Figure 2. (The nonlinear dependence on time of irradiation was due to change of light absorption. With initial absorbance below 0.25 the photolysis closely followed first-order kinetics.) The nitrogen balance (Table I, columns 8 and 9) also implies that no other nitrogen products are formed in considerable amounts. The following stoichiometric relations were found [Nz] = [NH,] = -(0.33 [Hz] N -(0.035
f
f
0.02) X A[NH20H]
0.002) X A[NHzOH]
The quantum yield of Nz was measured under conditions of total absorption: t$xz = 0.25 f 0.02, independent of the NHzOH concentration up to 1 M . In the presence of ethanol both and ~$-NH*OH decrease but ~ $ N Hand ~ C$H~hardly change; Nz and NHa still account for the nitrogen balance (Table I). The effect of alcohol increases with the ratio [CzHsOH]/ [NHZOH]; at ratios 5 0.1 no effect of alcohol could be discerned. I n acidic solutions the photolysis reveals new fea% tures: there is a pronounced decrease in t $ ~ relative to C$NH~ and NzO is produced. These results were observed a t pH values where almost all the hydroxylamine was in its acidic form; still a remarkable effect of pH The Journal of Physical Chemistry, Vol. 76, No. 8, 1973
1 , min
Figure 2. Yields of photolysis of degassed 2 mM NIlzOH solutions at pH 7.4, 23'.
was detected on lowering its value from 4.3 to 2.0 (Table I). At pH 2.0 NH3 and NzO are the major N-containing products, and no Hz could be detected. Owing to the high concentration of hydroxylamine a large error was involved in determining the yield of NH3 (10) Mass spectrometric analysis of the gaseous products showed traces of N20, the amount of which increased gradually on standing. This was probably due to thermal decomposition of NHzOH, which was carried with the gas, inside the mass spectrometer.
PHOTOLYSIS OF HYDROXYLAMINE IN AQUEOUS SOLUTION
20
183
v v
PULSE RADIOLYSIS
-
i
N
b
S '
15-
w
w ' 0
z
U m
%
u .
z
0
8U
8m
10-
*1
10
'Q
h
5.
210
220
230
240
s
250
1
2
2
( A15%)) but it appeared to be the same a t pH 4.0 and 2.0. B. Flash Photolysis. The flash photolysis of Nzsaturated solutions of hydroxylamine gives rise to a transient absorption below 300 nm. No other absorption was detected up to 740 nm. The same band was produced at pH 5.4 and 7.3, showing a well-defined peak 216 k 1 nm (Figure 3 ) . (The figure shows with A,, the measured spectrum and that corrected for stray was determined by the method of light effects. A,, midpoints. 11) The decay rate of the 230-nm absorption was studied at pH 7.6 and was found to obey a second-order law with 2k/e230= (1.0 It 0.3) X 105 cm 8ec-l. At pH 5.5 the rate constant was somewhat lower but still within the limit of error (Table 11). Figure 4 (insert) shows the pH dependence of initial transient absorption (Le., absorption extrapolated to zero time) at 230 nm. The measurements were limited to pH 54.2, where hydroxylamine is in effect totally in its acidic form and therefore its absorption should
I
3
4
5
I
B p H 3
6
7
8
4 9
PH
X, nm
Figure 3. Transient spectra produced by pulse radiolysis and flash photolysis of hydroxylamine solutions, corrected for stray light (solid curves, the spectrum produced by flash was normalized): 0, pulse radiolysis, 1 mM, NzO (20 mM), pH 7.6 and 5.4, 0.75 msec after pulse (each point is the average of six measurements, three at each pH, with a scatter of =!=5%); 0 , pulse radiolysis, 1 and 5 mM, N20 (20 mM), pH 2.1, 1 msec after pulse; V, flash photolysis, 10 (pH 5.4) and 2 mM (pH 7.3), Nz, 0.75 msec after flash. Dashed and dotted curves represent the spectra without correction for the pulse and flash experiments, respectively. . The limits of error for stray light are shown by lines, solid (pulse) and dashed (flash).
1 1; I..
FLASH PHOTOLYSIS
Figure 4. The pH dependence of initial transient absorption produced by pulse radiolysis and flash photolysis (insert) of hydroxylamine solutions. Pulse duration, 1.5 psec; 0, pulse radiolysis, 5 mM, Ar; V, pulse radiolysis, 10 miM, NsO (20 mM); 0,pulse radiolysis, 1 miM, Ar; 0 , pulse radiolysis, 1 and 5 mM, NsO (20 mM); A, flash photolysis, 20 mM, Nz.
Table I1 : Decay Kinetics of the Hydroxylamino Radical [NHzOH I, M
2X
PH
Method
7.3 5.4 7.6 5.5
Flash photolysis Flash photolysis Pulse radiolysis Pulse radiolysis
2k/azao, om sec-1
(1.O f 0 . 3 ) X (7.0 =t2 . 0 ) X (1.0 f 0.3) X (7.5 f 2.5) X
lo6 lo4 lo5 104
hardly vary with pH. Thus this figure might represent the pH variation of absorbance at constant yield of radicals. It shows a large change around pH 4 and levels off a t pH 5 2 , where the absorption was too OW for a kinetic study. The effect of ethanol was studied on solutions of 2 X A4 hydroxylamine, saturated with Nz at pH 6.5. With 1.4 X M ethanol the initial transient absorption (measured at 220 nm) was lowered but the decay kinetics hardly changed. Most of the absorption could be quenched on raising the alcohol concentration above 0.05 M , but a new transient spectrum emerged which mas weak, smeared, and extending to longer wavelengths. This absorption, probably being due to (11) G . Scheibe, Ber., 5 8 , 586 (1925). The Journal of Physical Chemistry, Vol. 76,No. 2, 1972
D. BEHAR, D. SHAPIRA, AND A, TREININ
184 the ethanol radical,12 could mask any residual absorption of the 216-nm transient. C. Pulse Radiolysis. The pulse radiolysis of hydroxylamine solutions gives rise to a transient absorption which is identical with that produced by flash photolysis. This is illustrated in Figure 3, which records the spectrum obtained with NzO-saturated solutions a t pH 5.4 and 2.6, and in Table I1 where some kinetic data are summarized. The extinction coefficient €230 = 1020 f 30 M-l cm-' at pH 5.5 was determined by dosimetry (EZM 1500 M-' cm-l) ; this leads to 2k = (7.7 2.4) X lo7 M-' sec-' for the rate constant of decay at pH 5.5. Somewhat different results have recently been reported by Simic and Hayon4 for the same transient: A,, 217 nm, emax = 2.5 X lo3 M-l cm-l, 2k/e240= (4.1 f 1.5) X lo5 cm sec-' and 2k = (4.5 f 2.0) X lo8 M-' sec-'. They also found some deviations from second-order kinetics, which we could not verify. The difference in pH cannot explain this discrepancy (see Table I1 and Figure 4), the reason for which is not clear. I n agreement witli recent w e assign this transient absorption to the radical NHOH, produced by reaction 1
-
OH
+ NH2OH +NHOH + HzO, kl = 9.5 X lo9 M-l sec-l
(1)
or by the slower reaction of OH with NH30H+. We could verify that the radical is neutral by not observing any distinct ionic effect on adding 0.1 M Na2S04to M hydroxylamine at pH 7.2. lll solutions of hyThe pulse radiolysis of 5.6 X droxylamine a t pH 5.5 produced the same amount of transient irrespective of whether they were satutated with NzO (2 X M) or Nz. Tinder these conditions about three-fourths of the solvated electrons are scavenged by XZOto produce OH radicals @(eaq- NzO)13 = 5.6 X 109 M-' sec-'; for the rate constant of eaqwith hydroxylamine see below). This implies that by attacking NHzOH, eaq-, and OH give rise to identical yields of NHOH. The following reactions can account for this result
+
+ NHzOH eaq- + NHsOHf eaq-
NH2
+ NHzOH
-+
NH2
+ H2O NH3 + NHOH
-+
----f
+ OH-
NH,
(2)
(3) (4)
and similar reactions with the species involved being in their acidic forms. Experimental evidence has recenty been provided for formation of the amino radical by reactions of eaq- with hydr~xylamine.~ Figure 4 shows the pH dependence of the initial transient absorption a t 230 nm. Most of the data were taken with 5 x 10-3 M hydroxylamine saturated with Ar, but other conditions were also employed to show that within the small dose variation, Figure 4 represents the change of absorption of a constant amount of hydroxylamino radicals (provided GOH and G, 4- GH The Journal of Physical Chemistry, Vol. 76, No. 8, 1978
are independent of pH in this pH region14). I n addition to the foregoing conclusion about the equivalent effects of eaq- and OH (compare the Ar and NzO experiments a t pH >5) it could be shown that OH and H radicals attack ?\TH30Hf to produce the same amount of the transient. Thus identical transient absorption M NH30Hf at were obtained with solutions of pH 2.1 irrespective of whether they were saturated with Ar (when -95% of eaq- are converted into H atoms) or with 2 X M NzO (when 60 and 37% of eaq- are converted into H and OH radicals, respectively) (Figure 4). Figure 4 is a typical titration curve, from which pK = 4.0 0.1 could be derived for the equilibrium NHzOH+
5
NHOH
kbf
+ H+
(5)
in good agreement with recent results (pK = 4.2 f 0.14). Under the conditions employed the decay of radicals could hardly disturb the acid-base equilibrium. 1O1O A I - l sec-l, i.e., k5 106 M-l Assuming that ks' sec-l, then under the conditions employed (concentraM ) the decay of NHzOH+ tion of radicals 5 2 X was always much slower than reaction 5. As for the back-reaction, up to pH 5 the decay of NHOH was much slower than protonation. At higher pH values, reactions with other Lewis acids (NH30H+, H2P04-, or water) should be considered, but their quantitative effect is not known. A high pH sensitivity of the absorbance was observed a t pH 25.5 (Figure 4), which might reflect changes in rate of protonation of NHOH by the various Lewis acids. (The pK values of NH3OH+ and HzP04- are 6.0 and 7.2, respectively.16) The flash photolysis results (Figure 4, insert) show similar featurcs but the titration curve is somewhat flatter. In this case considerable errors could be involved in measuring the low absorbance of the acidic form. The spectrum of this form obtained by pulse radiolysis is included in Figure 3. The rate constants of reactions 2 and 3 were investigated by following the decay of the eaq- absorption a t 600 nm in pulsed hydroxylamine solutions saturated with Ar (in presence of 1.3 x loe2M isopropyl alcohol to scavenge the OH radicals). Figure 5 shows the dependence of the pseudo-first-order rate constant on the hydroxylamine concentration at three pH values. No further change was observed on raising the pH from 7.7 to 8.2 (Figure 5). From these results we could dcrive IC2 = 6.6 X 10s M-l sec-', k3 = 1.0 X lolo M-l sec-I, and pK = 5.8 for the dissociation constant of NH30H+. The pK is in good agreement with availN
-
(12) M. Simic, P. Neta, and E. Hayon, J. Phys. Chem., 7 3 , 3794 (1969). (13) M. Anbar and P. Neta, Int. J . A d . Radiat. Isotopes, 18, 493 (1967). (14) G. Czapski, Advan. Chem. Ser., No. 81, 106 (1968). (15) "Handbook of Chemistry and Physics," 51st ed, Chemical Rubber Publishing Co., Cleveland, Ohio, 1970-1971.
PHOTOLYSIS OF HYDROXYLAMINE IN AQUEOUSSOLUTION
I
i's1 i 0 J 2
-
followed by reactions 4 and 7. Reaction 10 was introduced to account for the H2produced, with k9/klo 10 (since $ N H ~ / ~ 10). Mechanism I1 can be ruled out as the major one by considering the alcohol effect. I . Mechanism I1 predicts that not more than 50% of the yields of NHOH and Nz can be suppressed by a substance which can completely scavenge H atoms without reacting with NHOH. Our results indicate that ethanol behaves in this way since it affects neither the decay rate of NHOH (flash photolysis experiment) nor the yield of N2 when NH20H is present in excess (Table I). But much more than 50% of NHOH and Nz could be quenched by alcohol. 2. Under conditions where ethanol should effectively compete with NHzOH for H atoms to produce and (PNH~. This H2,'6 it exerts no distinct effect on is in complete disagreement with mechanism 11. Mechanism I is in agreement with the major features of our results. On complete scavenging of NH2 and OH by ethanol, NHOH and N2 should not bo produced, but the yield of NHa should not be affected by replacing reaction 4 by reaction 11.
-
pH 5.4
3
185
1
2
I
3
4
1
HYDROXYLAMINE, IO+M
Figure 5. Rseudo-first-order plots for the reaction of eaqwith hydroxylamine a t various pH values. (Full circle represents a measurement at pH 8.2.) The solutions were saturated with Ar and contained 13 mM isopropyl alcohol. Pulse duration: 0.2 rsec.
NH2
able data.15 The values of kz and ka are much higher than the recorded value IC < 2. X lo7 (at unspecified pH)lS but are close to that recently r e p ~ r t e d . ~
Discussion A , Photolysis of N H 2 0 H . Most of the photochemical experiments were carried out a t pH 7.4, where ca. 96% of the hydroxylamine is in its basic form. But even a t pH 5.4, where some flash experiments were conducted, NHzOH was the major light absorber. The stoichiometry of photolysis of NKzOH in aqueous solution is essentially identical with that of its Hgphotosensitized decomposition in the gas phase2 3NH2OH +NHs
+ 3Hz0 + Nz
The following mechanism (I) accounts for these results NHzOH
-% NH2 + OH
+ NHzOH +NHOH + HzO NHz + NHzOH +NHOH + NHs 2NHOH +Nz + 2H20 OH
(6)
(7)
The flash photolysis experiments provide direct evidence for the generation of NHOH. However, a second mechanism (11) should also be considered
+H H + NHzOH 4NHz + HzO NH20H h", NHOH
H
+ NHzOH
4NHOH
+Hz
(8)
(9)
(10)
4NH,
+ CzHiOH
(11)
To account for the small and constant yield of H2, a minor primary process should bo considered as occurring in parallel to reaction 6. The possibility that HZ results from some side reactions of the radicals (involved in mechanism I) can be ruled out, since in this case ( P H ~ should decrease on adding ethanol. A possible reaction that was already proposed for the gas photolysis2 involves internal recombination NHzOH h", H2
+ NOH (or HNO)
(12)
but NOH is expected1' to yield N20, which was not detected on photolyzing neutral solutions. Still reaction 12 cannot be discarded because under the condition employed some H E 0 could exist in solution as the dimer hyponitrous acid. It, is also possible that reaction 8 does take place t o a small extent and that reaction 10 is much faster than 9. I n this case, no change of $HEis expected to occur when ethanol competes with NH2OH on H atoms. The overall decomposition through this parallel mechanism is
(1) (4)
+ CzHsOH
2NHzOH +N2
+ Hz + 2H20
The ratio $N,/$H% = 9 f 1 can be accounted for by assuming that 89 and 11% of NH20H undergo reactions 6 and 8, respectively. However, a picture like this requires that on adding ethanol not more than 89% of (16) The rate constant for the reaction of H with ethanol is 2.5 X 107 M-1 sec-1 [P. Neta, G. R. Holdren, and R. H. Schuler, J . Phus. Chem., 75, 449 (1971)l. That for NHzOH is unknown but it is not likely to be considerably higher than 109. (17) 0. P. Strauss and H. E. Gunning, Trans. Faradau Soc., 60, 347 (1964).
The Journal of Physical Chemistry, Vol. 7 6 , No. 8, 1978
186
D. BEHAR,D. SHAPIRA, AND A. TREININ
NHOH should be quenched and that 4 N z should not fall below 0.54H,. We were unable to check these predictions because (a) the spectrum of the alcohol radical could mask any residual absorption of NHOH; (b) the ratio [CZH~OH]/[NHZOH] had to be below -200, so that with sufficient light absorbed by NH20H the absorption by alcohol was relatively small. Therefore no limiting value of $Nz could be measured. According to the proposed mechanism 4 N z equals the total quantum yield of primary dissociation. The excited state of NH20H involved is dissociative,2 and since 4~~ = 0.25 irrespective of [NH20H]up to 1 M , an efficient primary recombination should occur. The direct photodissociation of NHnOH to NH20 H, which is a major process in the gas phase,2 does not appear to take place in solution. We have no explanation for this and for the close resemblance between the direct photolysis in solution and Hg-photosensitized decomposition in the gas phase. B. The Photolysis of NHSOH+. Below pH 4.5 almost all the light was absorbed by NH30H+, and the photolysis desplayed new features. They do not .necessarily reflect a change in nature of the primary processes, but they must involve some modification of the secondary reactions. Thus no change in the nature of light absorbing molecules is produced on lowering the pH from 4.3 to 2.0, but the photochemistry is considerably affected (Table I). NHOH is the only radical involved which has its pK within k l in this region. (The pK of NH2 is 6.7 f O.z4). Therefore we believe that this pH effect reflects different chemical properties of the basic and acidic forms of the hydroxylamino radical. Simic and Hayon4 presented evidence that the basic form is indeed NHOH, but since there is esr evidence for the structure NH20 in acidic solutions18and in solid NH20H.HCl,19the acidic form was considered4 to exist in some equilibrium with NH20. I n several works dealing with the photolysis of gaseous hydroxylamine2 and its electrolysis in solutionsjZ0 NHOH and KHzO were assumed to yield KZand NzO, respectively, on their annihilation. If this assumption is correct, then the present work supports the view that NHOH is transformed to NH20 (or NHsO+) in acidic
+
The Journal of Physical Chemistry, Vol. 76, No. 8, 1973
solutions. At pH 52.0, when almost all the hydroxylamino radical is in acidic form, reaction 7 is probably replaced by reactions 13 and 142
+ NHzOH NzO + HzO
2NH20 +HNO 2HN0
4
(13) (14)
and the stoichiometry of photolysis is now 4NHzOH +2NH3
+ NzO + 3HzO
Considering the low accuracy of our results concerning this pH region (Table I), they are not in discord with the proposed mechanism. I n addition, the quenching of suggests that the minor primary reaction leading to HZformation (possibly reaction 8 or 12) does not occur with NH30H+. C. Analogies with the Photolysis o j H2O2. NHzOH and Hz02 are isoelectronic and bear some resemblance. The dissociation energies NHz--OH and HO-OH are relatively low (61 and 51 lrcal/mol, respectively*), and therefore it is reasonable to expect the rupture of these bonds as the primary processes. The mechanism proposed for the photolysis of H202 in aqueous solution is 2 1 HzOz +2 0 H
OH
+ HzOz +HOz + HzO 2H02 + HzOz + 02
(15) (16)
(17) Reaction 16 is analogous to reactions 1 and 4, and the radical HOz is isoelectronic with NHOH. The spectrum of H 0 2 (A, 230 nm, emax 1.2 X loa 1l-l cm-l) resembles that of NHOH and it decays even more slowly than NHOH.22 (18) (a) C. ,J. W. Gutch and W. A. Waters, J. Chem. Soc., 751 (1965); (b) J. Q.Adams, 8 . W. Niclcsic, and J. It. Thomas, J . Chem. Phys., 45, 654 (1965). (19) H. Ohigashi and Y . Kurita, J . Phys. SOC.Jap., 24, 654 (1968). (20) (a) A. D. Goolsby and D. T. Sawyer, J. Electroanal. Chem., 19, 405 (1968); (b) G. R. Itao and L. Meites, J. Phys. Chem., 70, 3620 (1966). (21) C. H. Bamford and It. P . Wayne in “Photochemistry and Reaction Kinetics,” Cambridge University Press, New York, N. Y., 1967, p 33. (22) D. Behar, G. Caapski, J. Rabani, L. M . Dorfman, and H. A . Schwara, J . Phys. Chem., 74, 3209 (1970).
