SOTES
5200
5210 5220 5230 5240 Wave length, A. Fig. 1.-The influence of perchloric acid and of sodium nitrate on the absorption spectrum of a 0.02 M solution of neodymium perchlorate a t 25": curves 1, 2, 3 and 4 represent solutions 0.002,4,9.8 and 11.8 M i n HClOn. respectively. Curve 5 represents a solution 4 M in NaN08.
Results and Discussion Absorption spectra of 0.02 M aqueous solutions of neodymium perchlorate were measured a t 25' for a number of solutions containing perchloric acid in concentrations from 0.002 to 11.8 M . Typical rssults for the band system around 5200 and 5750 A. are presented in Fig. 1and in Table I. TABLE I ACID O S THE h l O I A R 0.02 M AQUEOUSSOLUTIOSOF I ' E o o Y m u n i €'ERCHI,ORATE AT 25" Band system around 5750 A.
T I I E IXFLUEXCE O B PERCHLORlC
ABSORPTIVITIESe
(llC104)
OB A
0.002
2
4
G
9.8
8
10.8
11.8
A,
A. 5700 5710 5720 5730 5740 5752 5760 5770 5780 5790 5800 5810 5820 5830 5840 5850
(€
PI,
1.50 2.82 4.54 5.24 6.57 7.15 6.68 5.52 5.47 4 97 3.06 3.01 2.GO 2.20 1.82 1.47
4 8 11 12 13 13 8 0
5 2 0 -2 -2 0 0 0
6 lfi
20 22 25 26 16 12 9 4 -1 -3 -3 -1 0 0
13 25 28 33 37 40 20 20 17 G 2 -1 -1 2 2 2
-
60)
x io3
2.5
5.5
40
GR
65 91 87 108 117 127 131 112 112 105 104 97 83 15 13 65
40 50 53 60 47 42 41 21 21 17 18 18
80 84 74 110 03 124 172 221 230 225 235 212 200 186 160 130
108 105 77 120 97 124 213 340 358 300 422 388 345 312 275 228
The spectral changes produced by perchloric acid in concentrations up to about 6 M are similar to those produced by a decrease of temperature.'b Absorptivities a t and around the absorption maxima increase slightly with increasing concentration of the acid, decreasing eventually a t the long wave length wings of the absorption bands. Wave
Vol. 63
length shifts of the absorption maxima were not observed. Such spectral changes are of a type usually described4 as a medium effect. Variations of the refractive index of the solutions with the concentration of perchloric acid should also contribute to the observed effect.6 Experimental evidence is insufficient to decide which part, if any, of the observed spectral changes is due to complex formation. Direct evidence for complex formation is however provided by the pronounced effect of perchloric acid a t concentrations higher than 6 M . In certain spectral regions, particularly above 5760 A. (see Table I) absorptivities are now very sensitive to variations of the concentration of perchloric acid. In the 5200 A. region (see Fig. 1) absorptivities change now partly in opposite direction than at concentrations below 6 M . In 11.8 perchloric acid the absorption maximum a t 5218 A. is shifted by about 6 8.;that at 4272 8. by 3 A. toward longer wave lengths. As shown in Fig, 1, spectral changes produced by 11.8 M perchloric acid are similar to those produced by much lower concentrations of nitrates. Complex formation is rather well established in the latter case.' It might finally be interesting to note that the concentration dependence of the apparent molar refraction of lanthanum perchlorate seems to indicate interaction between the lanthanum and perchlorate ion at (Clod-) > 6 The spectral behavior of the Nd+3-perchloric acid system is analogous to that of rare earthschloride systems, recently reported by J ~ r g e n s e n . ~ In both cases direct evidence for complex formation is obtained only under conditions of very high activities of the solutes arid lorn activities of the solvent. 2,, 644 (1934); G. Kortum, Z. (6) See N. Q. Chako, TIIIS J O U R N A L phyazk. Chem., B33,243 (1936). (7) R. E. Connick and S. W. Meyer, J . A m . Chem. SOC.,7 3 , 1176 (1951). (8) J. E. Roberts and N. W. Silox, zbid.. 79, 17S9 (1957). (9) Chr. K. Jorgensen, KgZ. Danske Vzdenskab Selskab, Mal. /as. Medd., 80, No. 22 (1956); Acta Chem. Scand., 10, 1503 (1956).
SPECTROPHOTOMETRIC STUDY OF T H E IONIZATION OF IIYDIthZOIC ACID I N AQUEOUS SOLUTION' BY EUGENE A. BURNS A N D FRANCES D. CHANG Jet Propulsion Laboratorv, California Institate of Technologa, Pasadena 5, Cali/ornia Received November 6 , 1058
The ionization of hydrazoic acid in aqueous solution has been investigated, in the past, by several authors. 2--6 These researchers, through the use of conductivity, potentiometry and alkalinimetry, have found and reported ionization constants for hydrazoic acid in aqueous solution in the range 1.7 (1) This paper presents the results of one phase of research carried out a t the J e t Propulsion Laboratory. California Institute of Technology, under Contract No. DA-04-495-0rd 18, sponsored by the Department of the Army, Ordnance Corps. (2) C. 4 . West, J . Chem. Soc., 705 (1900). (3) W. S. Hughes, ibid., 491 (1928). (4) H. T. S. Britton and R. A. Robinson, Trans. Faradau SOC.,28, 531 (1932). ( 5 ) M. Quintin, Compt. rend., 210, 625 (1940). (6) N. Yui, Boll. Inst. Phys. Chem. Res., 20, 390 (1941).
;
August, 1959
NOTES
to 2.8 x 10-5 mole/liter. Among these investigators, only Quintinb and YuiBdetermined the thermodynamic constant by extrapolating their data to zero ioiiic strength; their values are 2.8 x and 1.91 X 10-6 mole/liter, respectively. The present study has been prompted by the observation of a significant absorption spectrum in the ultraviolet region, resulting from the presence of hydrazoic acid; the existence of this spectrum presents the opportunity to evaluate the ionization constant by an independent method.
were treated in the manner indicated in the folloming derivation. Let A represent the absorbance observed a t a given wave length for a solution containing sodium azide, perchloric acid aiid sodium perchlorate
Experimental Materials.-Reagent grade chemicals were used at all times. Procedure .-The concentration of the stock sodium azide solution was determined by potentiometric titration with standard hydrochloric acid, using glass and calomel electrodes in conjunction with a Beckman Model G pH meter. The end-point was evaluated by taking the second derivative of the titration curve. The coefficient of variation of this titration procedure was 0.16700. Stock sodium perchlorate solutions were standardized volumetrically by reduction with excess titanous chloride, using osmium tetroxide as a catalyst, followed by back-titration with ferric ammonium sulfate, with potsmiurn thiocyanate as the indicator .7 The experimental solutions were prepared by mixing aliquots of stock sodium azide, sodium perchlorate and perchloric acid, and then diluting to the desired volume. The solutions containing hydrazoic acid were used within 1 hr. after mixing; however, it wm observed spectrophotometrically that these solutions were stable for long periods (e.g., 2 wk.) if the container was well stoppered and was open for a minimum exposure time during transfer to the optical cells. Absorbance measurements were recorded with a Beckman RIodel DK-2 quartz photoelectric recording spectrophotometer. Matched 0.994-em. silica cells, equipped with ground-glass stoppers, were employed. All spectra were run against a blank solution of sodium perchlorate of the desired ionic strength. Measurements were made in an air-conditioned room whose temperature was controlled a t 21.7 f 0.6".
Resu!ts Data were obtained, in the region 310 to 220 mp, 011 the ultraviolet absorption of 0.0100 F sodium azide in the presence of varying amounts of perchloric acid (0, 0.00200, 0.00400, 0.00600, 0.00900, 0.0100, 0.0110, 0.0200, 0.0500 and 0.100 F) maintained a t a constant ionic strength. In the absence of perchloric acid, the absorption curve rose rapidly a t about 260 mp to a value of 1.0 a t about 250 mp. As the acid concentration was increased, the absorption curve began to rise sooner (300 mp), exhibiting a maximum a t about 260 mp and a minimum a t 232 mp, before attaining a value of 1.0 a t 221 mp. An isobestic point was observed, which is located at 254 mp and at 0.45 absorbance units. Perchloric acid concentrations greater than 0.0200 F produced no change in the spectra; hence, it was evident that the sodium azide had been completely converted to hydrazoic acid ( L e . , the solution contained 0.0100 M hydrazoic acid and 0.0100 M free perchloric acid). Inspection of the curves reveals that the minimum a t 232 mp is relatively sensitive to changes in acid concentration in the range 0.0090 to 0.0110 F, whereas the maximum at 260 mp is not. Hence, the region 228 to 242 mp is the optimum location for the spectrophotometric determination of the ionization constant. The data (7) E. A. Burns and R. F. Muraca (to be published).
1315
d = e~,(N3-)1
+ errNa("a)l
(1)
where (N3-) and (HN,) are the concentrations of azide and hydrazoic acid present in the solution, I is the path length of the optical cells, and e is the molar absorptivity of the species denoted by the subscript. Let (Hj0 and (N& equal the analytical concentrations of hydrogen ion and azide ion, respectively. Then A =
f?N,l(N3)0
f
(eHNs
- eNa)l(Hx;a)
(2)
The degree of dissociation of azide a is thus expressed by "3) a = - (=
(N3)O
[ . ~ / ( N ~ ) O-I [ ~ H N I- e N a
=
-
-e - eHNa
[A/(N3)011 eNt
m
(3)
When the degree of dissociation is known, the concentratioii ionization constant of hydrazoic acid k = (H)(N3)/(HN3)is calculated from the expression k =
(4)
The most accurate estimate of the ionization constant will be obtained when the analytical concentrations of azide and hydrogen ions are equal. Because of the molar absorptivities of azide ion and hydrazoic acid and the ionization constant of "3, TABLE I MOLARABSORPTIVITIESOF HYDRAZOIC ACID Wave length. lllp
228.0 230.0 232 0 234.0 2%. 0 238.0 240 0 242.0 250.0 260. o
270.0 280.0 290.0 300.0
ION -illolar
e Nz 518 483 445 411 36:3 327 282 232 8'3 18 3
..
.. ..
AND
AZIDE
absorptivity.-l./mole om. eHNa
33.0 28.2 26.2 20.4 27.4 20 1 31.2 33.5 41.2 46.0 39.3 24.7 12.1 4.1
the most accurate estimates of k may be obtained by fixing the analytical concentration of sodium azide a t an upper level of 0.0100 F and observing the resulting absorbance of solutions containing 0.0090 to 0.0110 F perchloric acid. The molar absorptivity of azide ion in the wave length region 228 to 242 mp was determined by averaging the values obtained from four different concentrations of azide ion (ie.,0.200, 0.400, 1.00 and 2.00 m F KaN3). The reproducibility of this procedure was good. The molar absorptivity of hydrazoic acid was evaluated from spectra obtained
NOTES
1316
TABLE I1 FOR HYDRAZOIC ACIDAT 22” AND IONIZATION CONSTANT Formal perchloric acid concn.
Wave length, mp
(Hh
242.0
0.0110 ,0100 .0090 ,0110 .0100 ,0090 ,0110 .OlOO ,0090 .0110 ,0100 ,0090 ,0110 .OlOO .0090 . 01 10 ,0100 .0110 .OlOO .OllO .OlOO Av .
240.0
238.0 236.0 234.0
232.0 230.0 228.0
a
Vol. 63 AT
VARIOUS VALUESOF IONIC STRENGT~
Ionization constant k
7
too=IIf
1.%O”M
0.
4.24 4.39 4.26 4.17 4.29 3.83 4.10 4.13 3.88 4.01 3.94 4.26 3.87 3.73 4.31 3.99 3.91 4.30 4.03 4.47 4.21 4.11 f 0.20”
3.99 4.84 4.73 4.24 5.13 4.35 4.33 4.75 4.31 4.58 3.44 4.89 4.83 4.22 4.67 5.28 4.10 4.38 4.49 4.87 5.10 4.60 f 0.35*
x
106
-
O.$OoM
o.l”oo=o A!!
0 . 0 ” O ~A t
4.24 3.98 3.98 3.97 4.02 3.84 4.09 4.31 4.50 3.85 4.10 4.30 4.20 4.36 4.47 4.70 4.61 4.86 4.00 4.46 3.91 4.23 f 0.29s
3.57 4.18 3.73 3.84 4.02 4.24 3.97 3.80 4.11 3.71 3.70 3.66 3.52 3.89 3.95 3.95 3.06 3.95 3.62 4.07 3.94 3.86 f 0.20”
3.32 3.52 3.37 3.00 3.01 3.28 3.59 2.93 3.28 3.14 3.25 3.04 3.02 3.39 3.38 3.51 3.27 3.64 3.37 3.32 3.27 3.28 f 0.20“
Standard deviation.
I
I
I
tained iii a similar manner. Table I lists the results of the determinations of molar absorptivities. The absorbance A of five series of solutions was recorded at 2 mfi intervals over the wave length region 228 to 242 mp, Each series of three solutions contained 0.0100 F NaN3, varying amounts of HC104 and sufficient NaC104 to maintain the desired constant ionic strength. The concentration constant le, which was calculated from eq. 3 and 4, is listed as a function of wave length, analytical concentration of perchloric acid and ionic strength in Table 11. Discussion The thermodynamic ionization constant for hydrazoic acid K was estimated, using the extended form of the Debye-Huckel equation as an npproximation of the activity coefficient product
I
In the case of hydrazoic acid, AZ x 2 i
=
2; a t 22’,
i
= 0.5064.8 Refining the function F by the expression
s(,)
0
1.0
I
II
1
1.2
I
1.3
PARAMETER A , (liters / mole)
Fig. 1.-Variance
I
Y
I.
sf as a function of the parameter A .
from solutions which contained 0,0100 F NaNa, 0.100 F HCIOl and which were maintained a t constant ionic strengths of 0.0300, 0.100, 0.200, 0.500 and 1.000 M . Molar absorptivities of azide and hydrazoic acid in the region 250 to 300 rnp were ob-
and estimating the parameter A , the function F was plotted us. p as suggested by Guggenheim and Schindler,g and the best straight line was evaluated by the method of least squares.lO The estimate of (8) H. 6. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” Reinhold Publ. Corp., New York. N. Y . , 2nd ed., 1950, p.567. (9) E. A. Guggenheim and T. D. Schindler, THISJOURNAL, 38, 543 (1934). (10) W. J. Youden. “Statistical Methods for Chemists,” John Wiley and Born, Inc., New Y o r k . N. Y., 1951, pp. 40-45.
NOTES
August, 1959
1317
variance F 2 of a single F measurement, compared with the least-squares line, was evaluated as a function of the parameter A and is plotted in Fig. 1. The minimum variance and, hence, the best fit to a straight line occurred a t A = 1.2; accordingly, the least-squares equation of the line is F = -4.622 0.221 f i . It follows that the thermodynamic ionization constant is 2.39 X 10-6 mole/liter, and 0 in eq. 6 is 0.221 liter/mole. The error in the constant K is estimated to be zt0.15 X mole/liter. TABLEI11 ACTIVITY COEFFICIENTS OF AZIDEION YNa-
0.727 0.002b 0.81 ,618 .86 .72 .565 .83 .68 ,519 ... ... 1.000 ,581 ... ... Data from ref. 11. Interpolated from data of ref. 11. 0.0300 ,1000 .200 ,500
a
Table 111 lists the values of the product of the activity coefficients YH + Y N ~ -observed in this work at 22'. Also listed are Kielland's values for the activity coefficient of hydrogen ion and for the calculated activity coefficient of azide ion.'l It should be pointed out that Kielland's data originated from work carried out a t 25O, whereas the present study was made a t 22'; these estimates do not include any correction for this discrepancy.
t,MINUTES.
Fig. 1.-Adsorption of CO on ZnO: O', Po = 51.4 cm. DBP; 25', Po = 49.8 cm. DBP; looo, Po = 59.5 cm. D B P ; 147O, Po = 42.3 cni. DBP; 20O0, Po = 53.8 cm. D B P ; 257", Po = 54.2 cm. DBP.
(11) J. Kielland, J. A m . Chem. Sac., 69, 1675 (1937).
a t 400°, so that the following run was made on a surface of different characteristics. At or below 100" a pressure of T H E RATES OF ADSORPTION OF WATER 10-6 mm. could be obtained a t the end of a run by evacuating the adsorption temperature for about 20 minutes. On AND CARBON MONOXIDE BY ZINC OXIDE at then raising the temperature to 400°, however, an appreciable amount of gas was desorbed: an additional 4 hours of BY MANFRED J. I>. Low A N D H.AUSTISTAYLOR evacuation were required to reduce the pressure to 10 -6 IVm. H . Nichols Laboratory, N e w York University, New I'ork 59, mm. At run temperatures of 147" and above, pumping New I'otk for about 8 minutes at the end of a run reduced the pressuie t o 10-6 mm. On then raising the temperature to 400 , Received November 7, 1968 more gas was evolved for some 4 hours. The chemisorption of CO on ZnO :It low temThese results thus directly contradict the statements of perat,ures is generally believed to be conipletely Garner, etoaal.,that the adsorption of CO is reversiblc below 100 . Rather, it appears that the adsorption is irrereversible mid ~ion-activated.~-~ Some rates of about versible and that more than one type of adsorption codesorption of CO have been m e a s ~ r e dbut , ~ with the exist over the entire temperature range covered. Such an exception of several experiments of Burwell and unstabIe system precludes reproducibility and vitiates atTaylor4 no adsorption rate data appear to exist. tempts :It quantitative correlation,q Iietwecn kinetics and or pressurc? on tho snmc :Idsorbcnt, bccause a Similarly, although the adsorption of HzO on tempei'nture different surface is obtained after each degassing. HowZnO has been examined by Taylor and 8icl
