D. L. CASON AND H. 11. NEUMANN
1822
boron. I t is proposed that a similar steric interaction toward the N202 group prevents the formation of stable adducts with these amines. Small amounts of the alkyl ammonium nitrites are obtained instead. Some definite trends in the stability of the aminenitric oxide products have been noted. The ammonium salts of the straight chain primary amines show a stepwise increase in stability from methylamine through n-butylamine. The n-amylarnine product is comparable in stability to that of methylamine. Isopropylamine and cyclohexylamine form compounds which are more stable than those of the straight chain primary amines. The stability difference of the sec-butylamine and isopropylamine product is surprising. The nitric oxide products of symmetrical, straight chain secondary amines exhibit a steady increase in stability from the diethylamine through the di-n-hexylamine products ; relatively, the latter two members of the series are very stable. Two branched-chain analogs, diisopropylamine and di[COA’TRIBUTIOX FROM
THE SCHOOL O F
Yol. 53
sec-butylamine do not form product but give very low yields of the alkylammonium nitrites. These observations on “stabilities” are purely qualitative and are based upon visual observation of the ease of decomposition of the material. I t has not been possible to correlate all of the “stabilities” described above with any single property of the amine donor. The basicity of the amine, crystal lattice effects, steric effects and volatility of the amine are all important considerations. With few exceptions, the stability of the products parallel the basicity of the amines as measured by the heat of formation of the triinethylboron adducts. There are some exceptions to this general :ule in instances where one or more of the above factors become controlling. Acknowledgments.-The authors would like to express their thanks to the Department of the Army Ordnance Corps for the financial support of the research under Contract KO. DX-11-022-ORD 2772.
CHEMISTRY, GEORGIA IXSTITUTE
OF
TECHNOLOGY, ATLANTA, GEORGIA]
Stability of the Chloro-complexes of Iodine in Aqueous Solution1 BY D. L. CASONAND H. AT. KEUMAKN~ RECEIVEDACGLXT26, 1960 Spectrophotometric measurements have been used t o study the stability in aqueous solution of the chloro-complexes of iodine in the (O), ( I ) and (111) oxidation states. The ion 12C1- is characterized by absorption maxima a t 248 and 437 mG, Qualitative observations were made on the conditions necessirj- for disproportionation of 12C1- to I- and IC12- and for the air oxidation of IzC1- to IC12-. The dissociation constant for IyC1- was determined bh- a spectrophotometric method, and the value so obtained was compared with the varietJ- of values in the literature. The most probable value is 0.60. The ion IC12- is characterized by absorption maxima a t 224 and 313 nip, and the spectral data on its dissociation are consistent with the work of Faull. Spectral evidence for IC1,- could not be found, either in solutions prepared from IC1, or from KI, K I 0 3and HCl. Solutions equivalent to iodine (111) in terms of oxidation were found to be equimolar mixtures of ICln- and Io3-.
The existence of chloro-complexes of iodine in dissolve salts. I t was then thoroughly dried and purified by the (0), (I) and (111) oxidation states has been sublimation. Iodine monochloride mas made by treating the purified well established. However, there are a number of iodine with condensed chlorine according to the method of conflicting statements in the literature as to the Cornog and Karges.4 I t was purified by two or three restability of these complexes in aqueous solution. crystallizations. .4 modification of the method of Booth and RIorris5 was The purpose of this work is to determine the condifor the preparation of iodine trichloride. Sinall tions under which these complexes exist in aqueous used amounts of liquid IC1 were added, a drop or two a t a time, solution. I t was hoped that this evidence would t o condensed chlorine. After all the IC1 was added, most of aid in interpreting the behavior of astatine in the excess chlorine was allowed t o evaporate froin the solid chloride solutions. Evidence for the existence of IC1,. IVhen solutions of IC13 were desired some IC1, would be shaken out onto a piece of paper, allowed a chloro-complex of astatine has been p r e ~ e n t e d , ~tcrystals o stand for a few seconds to let the excess CI? escape and but i t has been difficult to make arguments about then added t o the solvent without further delay. the oxidation state without further information d l l other chemicals were commercial materials of reagent grade. The concentrations of stock solutions were deterabout the behavior of iodine. The work reported here has established the nature mined by the standard iodometric method. Spectral Measurements.-The absorption spectra were and amount of the iodine species in chloride solu- determined with either a Beckman Model DU spectrophotions where the iodine concentration is in the range tometer or a Beckman Model DE(-1 spectrophotometer. lov5 to M . In addition to the experimental Most of the measurements were made with matched, fused work, which has principally involved spectrophoto- silica cells having a light path of 1.000 =k 0.003 cni. -4 water jacket was used t o keep the cells at a constant ternmetric measurements, the earlier data in the liter- perature of 25.0 & 0.5’. Even though the cells were ature has been evaluated. cleaned repeatedly, fluctuations in readings occasiorially appeared due to impurities in the cells. This difficulty with Experimental iodine solutions has been observed and discussed by other Materials.--;\ sample of commercial iodine was ground into powder arid treated with successive portions of w‘iter t o
workers.6
(1) Taken in part from a thesis submitted by Dennis Lamar Cason in partial fulfillment of t h e requirements f or t h e M S. degree, June 1959. (2) Requests for reprints should be addressed t o this author. (3) H. M. Xeumann, J . Inorg. Xuclear C h e m . , 4, 349 (1987).
(4) J. Cornog and R . A. Karges, “Inorganic Syntheses,” Vol. I , XlcGraw-Hill Book Co., S e w York, N. Y . , 1939, p. 165. ( 5 ) H. S. Booth and W. C. Morris, ibid., p . 167. (6) J. H.Wolfenden, Anal. Chein., 29, 1098 (1957).
CHLORO-COMPLEXES OF IODINE
April 20, 1961
The data are expressed in terms of the absorbancy, A , of a particular solution or in terms of the molar absorbancy index, a, of a particular species. For a solution containing a single absorbing species these will be related by the equation, A = a c 1, where c is the molar concentration of the species and I is the cell thickness in cm. Dissociation Constant Experiments .-For the equilibrium 12C1-(aq.) = I?(aq.) C1- (aq.) the dissociation constant is given by
+
if activity coefficients are ignored or by
if they are not. For absorption at a wave length where LC1- arid IY are the only absorbing species
+
A = (aI:cl-[InCl-l u1,~Inl)Z T o determine the dissociation constant, a series of experiments is performed in which the stoichiometric concentration of iodine is fixed at GO, ; . e . , co = [I21 [I?Cl-], but in which [Cl-] is varied. Letting AI, = a , c d and A I ~ c ~=UI,CI- col, appropriate substitution and manipulation give
+
If AI, is known and F can be estimated reasonably well, A can be plotted versus ( A - A I , ) / F / C ~ - ] . The slope of such a plot will be - K,, and the intercept 011 the A-axis will be illZcI-. Thus, in addition t o providing a determination of K,, the spectrum of I1C1- can be reconstructed. The method used in estimating F will be discussed in the section on the solubility of I?. One series of measurements was made in the visible region, using the DU spectrophotometer and the one cm. cells. I n this series the slit width was kept constant a t any one wave length so that the readings would be comparable. A second series was made in the ultraviolet region, using the DK spectrophotometer and 1 cm. cells containing 0.9 cm. silica plugs to reduce the light path to 0.1 cm. I n the latter series dust impurities were particularly troublesome. After a thorough initial cleansing the cells and plugs were kept immersed in a 6 M HC1 solution when not in use. Rubber gloves were worn during handling operations. The spectra were recorded immediately after filling the cells and then periodically during the next 30 minutes. The data mere considered invalid if the spectrum changed by more than a few per cent. during this time interval.
1823
are those of the mixture. If the solution is allowed to stand the absorption maximum a t about 440 mp drops and a new peak rises near 340 mp. The titer of the solution, as determined by iodometric titration, changes a t the same time; when the 440 mp peak has entirely disappeared, the titer has doubled its original value. Trial runs with various concentrations of HC1 in the range 1-12 M showed that the rate of disappearance of IZwas strongly dependent on the acidity, and incidentally also on the volume of air in the closed flask. These facts suggest that the reaction is air oxidation to IClZ-. From the practical point of view i t was observed that the spectrum of a 3 M HC1 solution only changes 10% over a period of 11 days, so that it is feasible to make stock solutions a t this or lower acidities. Both the ultraviolet and visible peaks of solutions of Iz in 3 M HC1 were found to deviate from Beer’s law a t sufficiently low iodine concentrations. The visible peak adhered satisfactorily to the law for iodine concentrations greater than formal. Solutions about formal in Iz have a peak in the ultraviolet a t 248 mp (IZCl-) with a hump on the short wave length side suggesting the presence of some IC12-. This spectrum may be observed within minutes after preparation of the solution, and it does not change with time. The spectrum is the same in both 3 hl HC1 and 3 M KC1. Llrhen the iodine concentration is reduced further, the peak appears a t 224 mp with a hump a t 248 mp. These effects provide evidence of the disproportionation reaction 12c1-
+ c1- = 1c12- + 1-
Dissociation Constant of IzC1- by Spectral Method.-Several conflicting values of the dissociation constant of IZC1- have been reported, so t h a t an independent measurement and a re-evaluation of previous results seemed in order. The principle of the spectral method has been outlined in the experimental section. Measurements were made in the region of both the ultraviolet and visible Results Spectral Characteristics.-Table I records the peaks, using HCl solutions in the concentration values of Xmax and amaxdetermined for species of range 0.1-6 M. The iodine concentrations were interest in this study. Comparison with values compromise values; large enough to keep the previously reported in the literature have been extent of disproportionation small but small enough to give reasonable absorbancy readings made where possible. with the cells available. Figure 1 shows the TABLE I spectra obtained in the visible region for a series Species Amax amax Ref. of solutions in which the stoichiometric iodine 1c12343 243 This work concentration, Co, is 5.33 X formal and the 343 275 light path is 1 cm. Determination of the dissocia224 47,000 This work tion constant from data in this region has the ad12c1437 1,100 This work vantage t h a t species which might be present as b .... 247 minor constituents, such as I-, 13-, IC1,- 01 HOI, 248 55,000 This work do not absorb strongly in this region; but has the I2 4 60 728 This work disadvantage that the spectra of TzC1- and Iz 460 746 c overlap to a considerable extent. Figure 2 shows 2 70 121 e the spectra obtained in the ultraviolet region for a Ref. 33. * L. I. Katzin, J . Chem. Phys., 20, 1165 a series of solutions in which Co is 2.60 x A. D . Awtry and R . E. Connick, J . A m . Chem. formal and the light path is 0.1 cm. In this (1952). Soc., 73, 1842(1961). region the absorption by IzC1- is considerably Iodine(0) .-When solid 1 2 is dissolved in hydro- greater than that by Iz,which is a favorable circhloric acid thei odinei n solutioni s present ini- cumstance; but the possible minor constituents tially as 1 2 and IZCl-, and the spectral characteristics absorb strongly. 0
1s2-4
D. L. CASONAND H. M. NEUXANN
1
I
I
1
:
I
Vol. 83
,
390
430 470 510 Wave length, mp. Fig. 1 .-Spectra in the visible region. Curves, from top t o bottom a t 4-20 mp, are: reconstructed spectrum of 12Cl-, q x c t r a in 5.7, 3.07, 1.23, 0.75, 0.31, 0.20, 0.098 dl HCI :tnd in 0.12 HCIOI.
210
230 250 2T0 29il Wave length, inp. Fig. 9.-Spectra in t h e ultraviolet region. Curves, from top t o bottom a t 250 mp, are: reconstructed spectrum of I & - , spectra in 2.90, 2.45, 2.02, 1.55, 1.12, 0.94. 0.75, 0.57, 0.39, 0.20 M HCI and in 0.10 M HClOI.
01
02
03
04
05
( A - AI,)/F[C!-] Fig 3 --Determination of the dissociation constant for 11'21- from d a t a in the visible region, d a t a for wave lengths of 440, 420, 410 and 400 mp
of -4 vs ( A - .41,)/F[Cl-l for selected wavelengths. The slope which best satisfies all of the data is - 0 60 The intercepts from plots such as these give the reconstructed spectrum of 1&1shown in Fig. 1 The spectra of the HClOd solutions showed greater variability among themselves in the ultraviolet region, but this was chiefly in the region around 260 mp where the absorption is small compared to that of 12Cl-. This is apparent in Fig 3 where the spectrum in 0 1 M HC104, appropriately corrected for iodine concentration, is shown It is clear that there is no real isosbestic point; the minor constituents I- and IC12- are contributing to the observed spectra in the HC1 solutions Figure 4 shows the plots of .1 vs (.IAl~,)/FIC1-]from these spectra In this case the slope is again - 0 GO. I n both cases the slope is not very sensitive to the values chosen for F, and the same slope is obtained by taking F as unity
In order to provide information on how the spec_. 1 1 trum of iodine is affected by acidity, the spectra of 1.1,5 x 10-' formal solutions of Ia in 0.1-1 h f HC104 were examined, using 1 mi. cells. lower I2 concentration was used in the HC104 experiments than in the HCl experiments to make the spectra more comparable from the point of view of possible disproportionation, since [I2]is less than co 08 -i in the HC1 solutions. In the visible region the spectra of the HC10, solutions were nearly identi06 cal. There was, however, a small increase (loyo 01 a t most) in absorbance with increasing acidity in the region of 520 mp. This observation sug!I 2 gests that a small fraction of the I2 may exist in I a non-solvated condition' in HC104 solutions. ~ _ _~~. . -2. .. L i The spectra of these solutions are sufficiently 0.2 0.6 1 .o 1.4 1.8 alike so that any one of them could be used t o ( A - .4r,)/F[C1 -1. give III, for the determination of the dissociation constant. The curve shown in Fig. 1 is for a 0.12 Fig. I.---Determinati(,n of the dissociation constant for *lI HC104 solution, corrected for the difference in IIC1- from d a t a in the ultraviolet region. Data for wave Til. The presence of an isosbestic point a t 470 m p Ipngths of 950, 2;i5, 2AO, 265 and 270 nip. is strong evidence for only two major forms of The principal limitation of this method is the iodine in these solutions. In Fig. 3 are shown plots difficulty in attaining conditions where only I? ( 7 ) J . G. Bower and R. I,, Scutt, J . A m . Chpriz. Soc., 75, 3.583 and 12C1- are present. However, the value of , I!KG) ~
_-L.
--_L
CHLORO-COMPLEXES OF IODINE
April 20, l O G l
1S"5
K,, 0.60, is in good agreement with the results obtained by Ray and SLirkar8 in distribution experiments. They obtained an average value of 0.62 from a variety of experiments using the nonaqueous solvents CS2, CC14 and CHC13, and HC1 solutions 0.06-1 J1 in concentration. Solubility of I, in Chloride Solutions.-The solubility of iodine in aqueous solutions of many chlorides has been determined by a variety of workers.8-19 Several of the workers used their own data to calculate the dissociation constant of I,Cl-. Among the values which have been stated explicitly or which can be obtained by simple cmversions are 200 250
0.48(
30" 35" 45O
.(i3 7;
~ ~ 1 1 1 3
.BO (HC1),S0.55 (HCI)," 0.81 ( H C P . 6 8 (?;aCli,14 0 . 6 1 ( S a C l i , " .63 (KCl),140.63 (KCl),Is 0 . 6 3 (KC1)18 (~c1,13 (~c1!19
. i 2 (HC1j19 Although a value of about 0.6 is strongly suggested a t 25', certain of the original data have subsequently been re-examined and re-calculated to give other values. For example, Sidgwick20 gives 0.48and Korenman21 0.29. These last two values, or values of A F o calculated from them, have subsequently appeared in articles of a review nature.22-'4 Considering the large amount of data available, the solubility measurements should yield a good value for the dissociation constant. For this purpose the solubilities reported were first all reduced t-, the basis of molar concentration for comparison of consistency. Density data were necessary for some of these conversions; the densities of the chloride solutions were taken frcm the International Critical Tables, and the assumption was made that the presence of I2 did not change the density. There is enough overlapping of data a t 25' from references S, 10, 11, 14, 13 and 1s to canfirm the reliability of their data. Certain of these data are shown in Fig. 5. The data of LVinther,I3 Ray17 and Sandvedi6 a t temperatures other than 25' are not inconsistent with these, but the extent to which they can be compared is limited. ( 8 ) P. R a y and P. V. Sarkar, J . Cheiii. Soc., 1449 (1922). 18) E . Oliveri-Mandala and A . Angenica, Gazz. c h i m ; i d . , 60, ( I ) , 2 7 3 (1920). (10) J. S. Carter, J . Chem. SOC.,28!il (19253. (11) J . S Carter a n d C . K. Hoskins, i b i d . , 380 11929). (12) W. Herz and I;. Hiebenthal. %. anorg. allgpm. Cheiiz., 177, 303 (1929!. (13) C . n'inther, Z . p h y s i k . C ~ Z ~3B, K , 299 (1929). (14) A . v. Kiss and A. Urmanczy, Z . aizorg. n l l g e i n . Ckeiiz.. 202, 172 1931). ! H . M ,Dawson and E . Spivey, J . Chenz. S O L . 1838 , (19132). ) K . S a n d r e d , Kgl. .\-orshe 17tderiskab. Selskabs Fovh., 6B, l(i3 (19.33). 17) S. K . Kay, J . l i z d i a i i C h e m SOL.,11, 11.5 (1934). (18) V. K . La hIer a n d M. W.Lewinsohn, J . P h y s . Chein., 38, 171 (19 3 4 ) . i l 9 ) F. H . Lee and H. C. Chen, .\-ankiiz~ J . , 7 , 3.53 (1987). \\'e are iiidebted to Dr. W. F. Linke f r r pruviding B gh3tu-copy of this article. ( 2 0 ) N . V . Sidgwick, "The Chemical Elements and Their Corni)iunds," Vol. 11, Oxford University Press, 19.50, p. 1193. (21) I . 11 Korenrnan, J . Gziz C h e m ( C . S S . R . ) . 17, 1C08 (1947) (22) R . L Scott, J . A m Chfiiz. Soc., 7 6 , 1551 (19.5':). I 2 3 ) R . C. Rrasted, "Comprehensive Inorganic Chemistry," Vol. 1 1 1 , D. Van h-ostrand Co.. 1934, p 185 21!i. 124) J . W. Xiellor, "A Comprehensive Treatise on Inorganic and Theoretical Chemistry," Supplement 11. IV31i. I) 8:37.
