R. M. WALLACE AXD E. K. DUKES
2094
Vol. 65
A SPECTROPHOTOMETRIC STUDY OF THE REACTIOS BETTVEES FERRIC ION AiVD HYDRAZOIC ACID' BY R. AT. WALLACE AND E. K. DVKES Savannah River Laboratory, E. I . du Pont de Nemours dl. Co., d i k e n , South Carolina Recemed July I O , 1061
A spectrophotometric study of the reaction between F e + + +and HK3 demonstrated that the complex FeNa+ T was formed. The molar extinction coefficient of FeNa++ was 3.68 X l o 3 a t the absorption peak of 460 mp. Equilibrium constants for the reaction F e + + + HN3 e FeN3++ H + were measured under a variety of conditions. The equilibrium constant a t 25", extrapolated to infinite dilution, was 1.71; the AH of the reaction was 2010 & 100 cal. Anomalous results for the equilibrium constant a t very low acid concentrations were attributed to the hydrolysis of F e + + +to FeOH++.
+
+
Introduction The reaction between aqueous hydrazoic acid and ferric ion to produce an intensely red-colored solution has been known for many years.2 Early investigators assumed that the color was due to the compound Fe(N&; however, Ricca3 has presented evidence to indicate that the reaction involved is
+ HN,
Fe+++
FeN3++
+ H+
(1)
Ricca's conclusion was based on the observations that the color faded upon acidification and aeration, that the color migrated toward the cathode in an electrolysis cell, and that the differences between the observed and calculated conductivities of solutions prepared by adding hydrazoic acid to ferric chloride could best be explained by the release of one hydrogen ion for every ferric ion undergoing reaction. Although the reaction has been used as the basis for a spectrophotometric analysis for hydrazoic acid,4 no systematic investigation of the reaction has been made. The purpose of the present investigation was to verify reaction 1 spectrophotometrically and to determine the equilibrium constant for the reaction. Experimental Materials.-Solutions of ferric nitrate, sodium azide, nitric acid, sodium nitrate, perchloric acid and sodium perchlorate were prepared from reagent-grade chemicals. Solutions of sodium azide were standardized by adding ceric ion to an aliquot of the solution and titrating the excess ceric ion with ferrous ion. Solutions of ferric nitrate were standardized by reducing the ferric ion to ferrous ion and measuring the absorbance of the ferrous o-phenanthroline complex. Apparatus.-Absorbance measurements were made with either :t Beckman DU spectrophotometer and matched 1.0-cm. Corex cells or with a Cary recording spectrophotometer and matched 2.5-em. Corex cells. The desired temperatures were maintained by using jacketed cells in conjunction with a thermostatically controlled water-bath. A Beckman Model H2 pH meter was used for pH measurements. Procedure.-Aliquots of standard solutions of sodium azide and ferric nitrate were added to a 10-ml. volumetric flask that already contained the other reagents such as nitric arid. After the solution was diluted to 10 ml. with water, :in aliquot was withdrawn for acid determination and the remainder of the solution was used for absorbance measurements. For measurements a t 0.01 Af H + , all reagents were added to a small beaker and diluted to about 9 ml. The solution was adjusted to pH 2 with dilute KaOH
_
_
_
~
(1) The information contained in this article was developed during t h e course of work under contract AT(07-2)-1 with t h e E. S. Stomlo Energy Commission. (2) Curtius and Rlssom, J. prakt. Chem., 118, Series 2, 261 (1898). (3) B. Riccs, Gazz. chim. %tal.,7S, 71 (1945). (4) C. E. Roberson and C. M. Austin, Anal. Chem., 89, 854 (1957).
or HNO3 and then \vas transferred to a 10-ml. volumetric flask for final dilution. Sbsorbance measurements m r e made against a reference of the same composition as the sample but containing no hydrazoic acid. When it was necessary to have large concentrations of H S 3present a t low acidities, as was the case in the determination of the extinction coefficient of FeK3++, the following procedure was used. An aqueous solution of HS3 (about 0.5 M) was prepared by passing a solution of YaN3 through a column filled with "Dowex" 50 cation exchange r e m in the acid form. The solution was analyzed for hydrazoic acid by a method described previously.6 Aliquots of the pure H E 3 solution, HNOI, Fe(NO&, and water then were mixed to obtain the desired final concentrations.
Calculations The equilibrium constant for reaction 1 can be expressed as follows, if activities are replaced by concentrations.
where the quantities in parentheses are the equilibrium concentrations of the respectire components. Since the ferriazide complex, E'eN3+T, is the only absorbing species in the region of the spectrum studied, and since the equilibrium concentrations of the ferric ion and the hydrazoic acid will be the concentrations of these components added less the amounts consumed in forming the complex, equation 2 becomes 4H-) K, = (3) e
[(Fe),
- A][(HS3?,
-
41
where A is the absorbance for a 1-cm. light path, E is the molar extinction coefficient of the ferriazide complex, and (Fe), and (HS3)nare the respective concentrations of the ferric ion and hydrazoic acid added to the solution. If the hydrazoic acid is present in large excess so that its concentration remains essentially unchanged by the formation of the complex, but an appreciable fraction of the ferric ion is consumed in forming the complex, equation 3 can be rearranged to give
h plot of (Fe),/A us. l/(HN3)&at constant acidity should therefore give a straight line with intercept l / e , permitting the determination of the extinction coefficient. (5) E. K. Dukes and R. 11.Wallace, .Anal. Chem., 33, 242 (19G1).
Nov. , 196:I
REACTION BETWEEN FERRIC ION AND HYDRAZOIC ACID
Results Application of the method of continuous variations6 to the system ferric nitrate-hydrazoic acid in nitric acid demonstrated that one mole of ferric ion and one mole of hydrazoic acid are involved in the formation of the colored complex. Figure 1 shows plots of absorbance z's. the mole fraction of HN3 for solutions in which the sum of the concentrations of Fe+++ and HK3 as constant. The results at 7 ilf HXOy were obtained with a total concentration of 0.16 Jrl (Fe+++ plus HK3) and those at 0.01 ilrl HNO, at a total coiiccntration of 4.0 X Af (Fe+++plus Hn-3). In each case the maximum occurred a t mole fraction 0.5, indicating the formation of 1:1 complex. Further evidence for the existence of a 1: 1 complex was the constaiicy of the quantity A/(Fe)&(HK,), at constant acidity and a t various concentrations of added Fe+++and HS3. The constancy of this quantity also shows that no significant amounts of Fe+++ or HN3were consumed in forming the complex in these measurements. The values of this quantity in 1 M HNO, for various concentrations of added Fe+++and HN3 are shown in Table I. The values are seen to be constant within experimental error, as would be expected for the proposed reaction. Similar measurements made in 5 , 6 and 7 M HK03 also gave constant values for A /( Fe)&(HX&. TABLE I
0.8
P
d E
;? 0.4
z
0.2
I
I
g. 0.6
2095
:Y 0.0010 0.0009
-
0.0008 0.0007 -
$ X0.0006 v
0.0005 0.0004
0.0003
VALLEE, OF d/(Fe)a(HN3)a IN 1 M "03 (Hh'da, M
(Fe)a. JM
A Absorbance
A (Fe)*(HNn)s
0 0096 0147 0147 0196 0245
0 0250 0125 0150 0125 0100
0.468 357 432 477 481
1950 1942 1959 1947 1963
0.0002 1 0
4.0
8.0 12.0 16.0 20.0 24.0 1/'("~)*. Fig. 2.-Determination of molar extinction coefficient of FeX3++ at 460 mp.
4.0
Determination of the Molar Extinction Coefficient of FeN3++.-The molar extinction coefficient of Fey3++was determined in 0.05 and 0.10 3.0 M HNO, by a method based on equation 4. In each determination, solutions were prepared containing the desired acid concentration, a small amount of ferric ion, and varying concentrations of I 0 hydrazoic acid. The absorbance of each of these 2.0 solutions was determined a t 460 mp, and the ratio X (Fe),/A was plotted us. the reciprocal of the hydrazoic acid concentration as shown in Fig. 2 . The resulting straight lines were extrapolated to 1/ (HS,)a equal zero, and the extinction coefficients 1.o were calculated from the intercepts. A least squares fit of the data gave extinction coefficients of 3.69 X lo3and 3.66 X lo3for the 0.05 and 0.10 M "03, respectively. The average value of 3.68 X lo3 n as used in subsequent calculations. 400 450 500 550 600 It was not possible to obtain reliable estimates of Wave length, nip. the extinction coefficients a t acid concentrations Fig. 3.--Absorption spectrum of FeXs +'. much greater than 0.1 &I; the higher acidities reduced the fraction of the iron present as the com- change when the acidity was doubled, it was asplex, and required a much greater extrapolation. sumed to be constant in all experiments. However, since the extinction coefficient did not The determination of the extinction coefficient a t one wave length together with absorbance measure(6) W C. Vosburgh and G R . Cooper, J . A n . Chem SOC, 63,437 ments at other wave lengths permitted the deter(1941). D
+
VI
R. 31. WALLACE AND E. K. DUKES
2096
T'ol. 63 TABLE I1 23' I N ASD iYaC1O4
EQUILIBRIUM COXSTASTS AT
KL
H+, M
0.50 M
0.01 .05 .IO .50
0.679 .765 ,718 ,723 ,735
Av.
0.2
0.4
0,6
0.8
1.0
(&if)"*.
Fig. 4.-Log 0.3
r-
K Ous. ( L W ) ~a' *t 23". I
--0.1
b
3 --0.2
c:
-- 0.3 -0.4
-- 0.5 --0.6
3.0 3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 l / T , OK. Fig. S.-Effect of temperature on K c .
mination of the absorption spectrum of FeSs++ shown in Fig. 3. The maximum in the curve occurs a t 465 rather than 460 mp, but the difference in the ex1,inction coefficients a t these two wave lengths is so small that 460 mp was used in all these studies. Effect of Acid Concentrate on Equilibrium Constant.--If equation 1 were valid, the equilibrium constants calculated from equation 3 should be independent of the hydrogen ion concentration, providing the ionic strength was held constant. Equilibyium constants therefore were determined in solutions containing mixtures of IIC104 and Sac104 in which the ionic strength was maintained constant at 0.50 and 1.00 d l and the acid concentration wa,3 varied from 0.01 to 0.30 AI. T i e results shown in Table I1 demonstratr that the equilibrium constant is independent of acidity over the range studied. The only significaiit discrepancies occurred in those nieasurements a t 0.01 M acid; these points have been disregarded in taking thc average, for reasons that will be discussed later. Although the acid concentration did not affect the equilibrium constant significantly, the ionic strength did affect it markedly. The average value of the equilibrium constant decreased from 0.735
HCl@
AIIXTURES O F
7
LOO
nr
0.545 .626 ,616 ,594 ,612
to 0.G12 when the ionic strength mas increased from 0.50 to 1.00 11.1. Equilibrium constants also mere determined in nitric acid solutions a t concentrations between 0.05 and 8.23 Jll and in perchloric acid solutions at concentrations between 0.05 and 1 ill in the absence of any added salts, so that the ionic strength was equal to the acid concentration. The results of these measurements are shown in Table 111. A comparison of the results obtained in nitric and perchloric acids a t the same ionic strength shows the equilibrium constants in the perchlorate system to be slightly higher than comparable ones in the nitrate system. The difference is not large but is greater than the experimental error and may be due to some specific interaction between the nitrate and ferric ions. Such an interaction would lower the actual ferric ion concentration without our having taken the decrease into account in our calculation of K , by equation 3 and would result in values of K , that are too small. EQUILIBRIUM
TABLE 111 AT 23" CONSTAXTS
IN
HX'Oz
ASD
fIC101
HNOI, M
Kc
H C I O ~,$r ,
Kc
0.0496 ,0992 ,198 ,496 ,744 ,992 2.05 3.05 4.05 5 13 6.15 7.13 8.23
1.210 1,059 0.913 .661 .596 ,550 ,411 ,340 ,301 ,257 ,272 ,259
0,0509 ,102 ,204 ,509 . 764 1.02
1.264 1.146 0.969 ,747 ,623 ,576
,
"jh
When log K c was plotted against t,he square root of the acid concentrations as shown in Fig. 4 for the data in Table 111, a reasonably good straight line was obtained for acid concentrations below 0.5 111 in the perchlorate system. Ext,rapolatioli of this line to zero ionic strength gave a value of 1.67 for the equilibrium constant at infinite dilution a t 23". A similar plot of the dat'a obtained in iiitric acid can be extrapolated linearly t o n value only slightly low-er t'han that obtained in perchloric acid. A smooth curve drawn through the points in the nitrate system, however, can be made t,o pass through the same point a t zero ionic strength as that found in the perchlorate system, The probable int'eraction between ferric and nitrate ions suggests that the linear extrapolation in the perchlorate system will yield the most reliable value for K , at infinite dilution.
REACTION BETWEEN FERRIC IONAND HYDRAZOIC ACID
xov., 1961
Although K , was independent of acidity a t lower ionic strengths, it varied slightly with acidity when the ionic strength was 5 M . Equilibrium constants of 0.370, 0.320 and 0.281 were obtained for solutions that were 1, 2.5 and 5.0 ill, respectively, in Hn'O3 to which sufficient S a N 0 3was added to maintain a constant ionic strength of 5.0 M . These variations are greater than the experimental errors. Effect of Temperature on Equilibrium Constant. Equilibrium constants were determined a t various temperatures in the range 20 to 58" in solutions which were 0.01, 1.0 and 5.0 A l in HXOS as well as solutions which were 1.0 and 2.5 iM in " 0 3 and contained sufficient l;aSOs to maintain a constant ionic strength of 5.0 M . The results of these measurements are shown in Fig. 5 in which log K , is plotted ayainst 1,'T. The straight lines that are obtained for all except the lowest acid concentration have negative slopes, which demonstrate that the reaction is endothermic. Enthalpies of reaction in the various media are shown in Table IV. The value obtained in 5.0 J I HK03 is significantly higher than the others, which indicates that the medium has an appreciable effect on AH a t higher acidities. An average of the first three values in Table IV, 2.01 ltO.10 kcal./mole, is probably the most reliable estimate of AH because of the above-mentioned effect of acidity. TABLE IV EWHALPY OF REACTION IN VARIOUS MEDIA Medium
1. 0 M 1.0 M 2 5M 5.0 M
HNOj 4 0 M SaN03 HXO3, 2 5 11.1 SaN03 HSOa
"OB,
AH, kcal. 'mole
$2.03 $1 89 +2.10
$2.42
Anomalous Behavior at 0.01 M "Ox.-The variation of K, with temperature in 0.01 M " 0 3 appears to be inconsistent with that found in the other solutions. The plot of log K , us. 1/T is not only curwd but it actually has a pocitive dope over most of the range of temperatures studied. The values of K,, however, were calculated with equation 3, which ignores the hydrolysis of ferric ion a t low acidities. The hydrolysis of ferric ion according to the reaction
+ H20
Fe+++
h',,
=
2097
+ Hi-
FeOH++ (FeOH++)(H+) ._ (Fe +++I
has been studied previously. Lamb and Jacques' report a value of 2.5 x for the equilibrium constant, Kh, a t %", while Arden8 found it to be 1.25 X lou3. Rabinowitch and Stockmayerg reported the enthalpy of reaction to be +l2.3 =t1 kcal./ mole. Using a value of 2.0 X 1 0 - ~for K h a t 25" and 12.3 kcal./mole for AH, we have calculated Kh a t various temperatures and corrected the apparent values of K , obtained in 0.01 M HKOi by multiplying them by 11 Kh/(H+)'. The corrected values are shown in Fig. 5. The line drawn through the corrected points now has a negative constant slope from which the enthalpy of the ferriazide reaction was found to be $2.38 kcal./mole. Although this result is not in perfect agreement with those obtained in solutions of higher acidity, it is sufficiently close to demonstrate that the apparent anomaly was due to the hydrolysis of the Fe+++ ion. It was because of this hydrolysis and the uncertainty involved in correcting for it that the results obtained previously in 0.01 21 KC104 were disregarded when averages were determined for K,. The following thermodynamic quantities for reaction 1 were estimated at 25" and infinite dilution from the value of K , a t infinite dilution at 23" and the enthalpy of the reaction.
+
K , = 1.71 AF2p = -318 cal. AH260 = 2010 & 100 c d . A&. = 7.8 lt 0.3 e.u.
From the value for K , given above and the dissociation constant for hydrazoic acid, 2.8 X reported by Quintin,'O the equilibrium constant for the following reaction was determined.
+ NB-
Fe+++
FeNaT+
(7) A. B. Lamb and -4.G. ,Jacques. J. Sm. Chem. Soc., 60, 1215 (1938). (8) T. V. Arden, J. Chem. Soc., 350 (1951). (9) E. Rabinowitch and R. H. Stockmayer, J. Am. Chem. Soc., 64, 335 (1942). (10) M. Quintin, Compt. rend.. 810, 625 (1940).
