NOTES
3722
VoI. 77
chloride ione even after periods of several weeks. Although it is well known that iodine inonochloride and iodine bromide conduct electric current in certain solvents and in the molten state,; very little quantitative work has, as yet, been done either upon their conductances in dilute solutions or on the mechanism of such conductance. I t was thought that such study may be of interest to the understanding of the behavior of iiiterh:tlogeii compounds in acetonitrile solutions.
0.04 3.0
I
I
,
3.4
3.2
r4 x Fig. 1.-Temperature
3.4
3.3 40
3.5
3.6
3
dependence of the U(1V) acid constant.
the entropy difference between U+4 and U+3 after correction for the entropy of hydrogen (Sf'~+l = -30 e.u., S?/,H,= 15.6 e . ~ . ~ the ) entropy change for reaction (1) readily can be accounted for. CHEMISTRY DIVISION OAKRIDGENATIONAL LABORATORY OAKRIDGE,TENNESSEE
Studies on the Chemistry of Halogens and of Polyhalides. IV. On the Behavior of Iodine and of Iodine Halides in Acetonitrile',' BY
I . P O P O V AND NORMAN E. RECEIVED MARCH28, 1956
ALEXANDER
sKELI,Y3
Spectrophotometric investigations of the solutions of iodine monochloride4 in acetonitrile have shown that this compound undergoes an ionic dissociation. 21c1 I + f 1c12It was observed, however, that the absorption spectrum of iodine chloride in acetonitrile changed with time. The ukraviolet peak a t 227 mM which is due t o the IC12- ion5 slowly increased in intensity, without, however, showing a quantitative conversion of iodine monochloride to the iododi(1) State (2) (3) (4) (5)
Abstracted in p a i t from t h e P h . D . thesis of Norman E. Skelly, University of Iowa, 1955. Previous paper of this series, THISJOURNAL, 76, 5309 (1951). Du P o n t Postgraduate Fellow, 1953-1954. R . E. Buckles and J. F. Mills, THISJOURNAL, 76, 4845 (1954) A. I. Popov a n d J. N. Jessup, i b i d . , 74, 6127 (1952).
Experimental Part Conductance Measurements.-The conductivity bridge, and the measurement conditions were described in a previou\ communication.? The temperature \$-as maintained at 25 & 0.(I03 . Spectrophotometric Measurements.--All spectropliotometric measurements w x e done on a Cary recording spectrophotometer model 11. Stoppered silica cells of path length 1.00 i 0.01 cin. were used. Measurement5 iverc' made a t the room temperature of approximately 3;". Solvents.-Acetonitrile obtained from the hlatheson, Coleman and Bell Co. was purified by the method already described.2 It is interesting t o note that different batches of the solvent seemed t o differ considerably in the amount of impurities present, and in some cases required several additional distillations over phosphorus pentoxide before solvent of sufficient purity was obtained. Iodine.-C.P. iodine ( J . T. Baker) was purified by sublimation from a mixture of calcium oxide and potassium iodide. Iodine Monochloride.-Iodine monochloride was prepared by the method of Cornog aud Karges.8 It was repeatedly purified by fractional crystallization in V Q C U O . The melting point was 9i.2' which is identical ivith the literature value. Iodine Cyanide.-Iodine cyanide was prepared and purified by the method of Bak and Hillebert.g The melting point in a sealed tube was 146-147", identical with the value reported. Iodine Bromide.-Iodine bromide was prepared by t h e addition of 88 g. (0.55 mole) of bromine t o 127 (0.50 mole) of iodine. The mixture was warmed on a water-bath until the iodine completely dissolved, When the mixture was cooled, solid iodine bromide crystallized out. The product was purified by repeatedly melting and cooling the crystals and discarding each time approximately 20yo of the liquid melt. This procedure was repeated until the same melting point was obtained on two successive recrystallizations. The melting point was found to be 41.5" as compared witli the value of 42" reported in the literature.10 Solutions.-Stock solutions were prepared by dissolving L: weighed amount of the solute in purified acetonitrile. The exact concentrations were determined by iodometric titrations. Aliquots of stock solutions were then diluted to required concentration. All the solutions were at first prepared in a dry-box and the absorption cells and the COIItluctance cells were likewise filled there. Since, as shall be described below, the absorption spectra as well as the specific conductance of the solutions showed a pronounced change with time, a special all-glass apparatus was designed which permitted the carryiug out of the purification of the solvent, preparation of solutions a t various known concentrations, and the measurement of the electrical conductance without any exposure to t h e atmosphere."
Results Iodine Monochloride.--Absorption spectra of iodine monochloride in acetonitrile have been re(6) T h e IClz- i u n has often been called t h e dichlotoiodide. However, since i t is produced b y t h e addition of t h e chloride ion t o iodine monochloride it seems more correct to call it iododichloride. ( 7 ) (a) L. Biuner a n d A . Galecki, Z . p h y s i k . C h e m . , 84, 513 ( I g l X ) , ( h ) Pi N. Greenwood a n d H . J . Emeleus. J . Clzcm. SOC., 987 (1960). (8)J. Cornog and R . A. Karper, "Inorganic Syntheses," Vol. I , McGrdw-Hill Book Co., Inc.. New York, N. Y., 1939, p. 165 (9) B. Bak and A. Hillebert, 0 7 2 . Synthescs, 32,29 (1959). (10) A. E. Gillam, Tians. F a r a d a y S O L . ,29, 1132 (1933). (11) A. I. Popov, iV,E. Skelly and R . Rygg, to be publishetl
NOTES
July 20, 1955 ported previously. Both maxima, the one a t 358 mp which is due to the iodine monochloride, and the one a t 227 mp, due to the IC12- ion showed increase in absorption with time (Table I). The rate of the reaction appears to be strongly dependent on the total concentration of iodine monochloride. In 4 X lop6 M solution of iodine monochloride after 17 days the reaction apparently goes to 95% completion (since the molar absorbancy index of IC12- a t 227 mp is 53400). On the other hand the 5 X M solution showed only approximately 13y0 conversion after 29 days. I t is interesting to note that in two cases the absorption spectra indicate the formation of the triiodide ion. TABLE I MOLARABSORBANCY INDICES AT 227 m p FOR IODINE MosoCHLORIDE IN ACETONITRILE AT 25" Time, days
-
Molar concn. X
500
0 334 4 514 7 720 10 Off scale 17 25 29 38 Appearance of
50
10
1080 1350 1482 1984 3480 3140 3496
4020 5020 7700 10900 a
1W
4
8
4762 6890 8938 12380
14450 15100 16890 21000
a
25000 25000 25000 25000
IS-.
The electrical conductances of iodine monochloride solutions in acetonitrile were found to be quite low, especially in the more dilute solutions. The molar conductances rose rapidly with dilution typifying the usual behavior of weak electrolytes. There was no possibility of obtaining the conductance a t infinite dilution by either the usual extrapolation methods, or by the method of Fuoss and Kraus. The measurements of the conductance were further complicated by the fact that it increased with time. Specific conductance of a 1.52 X lo-? M iodine monochloride solution prepared in the dry-box increased from 2.01 X ohm-' cm.-' to 11.36 X ohm-' cm.-' in 82.5 hours. The increase, however, was much slower when the conductance was measured in an enclosed apparatus mentioned above. A solution of 3.65 X M increased only from 1.18 X to 4.52 X 10-6 ohm-' cm.-'in 210 hours (Fig. 1). Iodine Bromide and Iodine.-The behavior of the iodine bromide and of iodine in acetonitrile were similar to that of iodine monochloride. Both compounds showed an increase in conductance with time and failed to follow Beer's law in the ultraviolet region of the spectruin. In this region iodine bromide and iodine had absorption spectra indicative of the dibroinoiodide and triiodide ions, respectively. The change in the conductance for iodine solutions was much less pronounced than that for the iodine halides. An increase of 0% in the conductance was observed for 0.158 hf iodine solution within a 24-hour period. The change in electrical conductance with time for iodine bromide is given in Fig. 1. Some values (12) R. M . Fuoss and C. A. Kraus,
THIS JOURNAL, 6 6 , 476
(1933).
3723
20
40
60
80 100 120 140 160 180 Time, hr. Fig. 1.-Variation of specific conductance with time: 1, iodine bromide solution 5.95 X AB; 2, iodine monochloride solution 1.54 X -21; 3, iodine monochloride solution in enclosed apparatus 3.65 X 1 0 F M . Divide ordinate values by 10 for curve 1.
for the molar conductances of iodine, iodine monochloride and iodine bromide are given in Table 11. The values were obtained in each case by the extrapolation of a series of measurements to zero time. TABLE I1 MOLARCONDUCTANCE OF IODINE A N D OF IODIKEHALIDES IN ACETONITRILE Iodine
C X 10%
15.80 7.74 4.10 1.90
0.50 0.10
Iodine mono;hloride
X Am
10'
Am
0.0349 3 . 1 4 0 . 3 2 5 ,0512 3.65 ,312 ,0920 3 . 7 6 ,237 ,132 .. ... ,245 ,, ,,, ,553 .. ...
Iodine bromide C X 10% Am
4.50 2.24 0.90 0.45 0.27
..
0.229 ,247 ,973 ,787 1.12
..
.
C X 101 44.8
.. ,.
.. , ,
..
Iodine cyanide Am
7 . 0 X 10-8
........ ,....... .,....., , , , , , , , . .... . . .
I t is interesting to note that although acetonitrile has a much higher dielectric constant than pyridine (36.5 VS. 12.3 a t 25O), yet the iodine exhibits a much smaller conductance in the former solvent. The difference in behavior is very likely due to the fact that pyridine is a much stronger electron donor than the acetonitrile and therefore in the Py.12 complex iodine is much more strongly polarized than in CH,CNnIz. Iodine Cyanide.-Iodine cyanide solutions in acetonitrile were extremely weak conductors, and a 0.448 M solution showed conductance of only 3.15 X ohm-' cm.-' which increased to 3.38 X ohm-' cm.-' after 81.6 hours. Because of the high concentration of the solute and the high resistance obtained even with the cells with smallest constant (0.0406) i t was thought impractical to carry out measurements a t greater dilutions since then the conductance of the solvent becomes an important factor. Discussion and Conclusions Iodine and iodine halides are notorious for undergoing some sort of a slow reaction in non-aqueous solvents, as was shown by both the electrical conductance and the spectrophotometric measure-
3724
Vol. 77
NOTES
ments, especially for the case of iodine-pyridine solutions. l 3 The reasons for the slow reactions have been attributed either to a gradual halogenation of the solvent13b,cor to a reaction with traces of moisture and/or other impurities.13e In a recent work of Buckles and Mills4 the authors report that they have found no evidence for a reaction of iodine monochloride with the solvent. It seems from the data obtained in this investigation that the evidence favors a slow reaction with the moisture because of the drastic change in the rate of increase of electrical conductance when the measurements were made in an enclosed system. It is evident, however, that a considerable amount of further work is needed to elucidate completely the behavior of interhalogen compounds in various solvents. Acknowledgment.-The authors gratefully acknowledge the support of a part of this project by the Research Corporation. They also wish to express their appreciation to Dr. Robert E. Buckles of this Laboratory for many helpful discussions on this problem.
Experimental Part Apparatus.-All spectrophotometric measurements were made on a Cary recording spectrophotometer, model 11, under conditions described in previous publications.lh Solvents.-Acetonitrile was obtained from the Matheson, Coleman and Bell Co., and purified b y methods already described.Ib Ethylene dichloride was obtained from the Union Carbide and Carbon Co. It was purified by sha.king it with several portions of sulfuric acid until the acid layer remained clear, washing it with copious amounts of water, followed b y washing with a dilute solution of sodium bicarbonate and again with water. Since ethylene dichloride forms a n azeotropic mixture with water, t h e latter was removed b y distillation and t h e last traces of moisture were further removed b y refluxing the solvent for several hours over phosphorus pentoxide and then distilling it twice from the same reagent. The boiling point of the solvent way 53.5" a t 760 mm. I t was found importaut to check the purity of ethylem dichloride by the following procedure. A small amount of bromine was added t o a couple of hundred milliliters of thc solvent and the amount of halogen was titrated iodometrieally in a n aliquot portion. T h e titration was repeated thc next day and if no change in bromine concentration was ohserved, the solvent was used for subsequent spectrophotometric work. Polyhalogen Complexes.-The following polyh~logcii complexes were prepared according t o the general directions of Chattatvay and Hoyle2: Tetramethylammonium di(13) (a) L. F. Audrieth and E. J. Birr, THISJOURNAL, 6 6 , 668 iodobromide, tetramethylammonium iododibromide, tetra11933); (b) G. Kortum and H. Wilski, 2 . p h y s i k . C h o n . , 202, 35 methylammonium iododichloride, tetranieth\-latnInoniuni (1933); ( c ) R. Zingaro, C . .VanderWerf I and J Kleinberg, THIS iodobromochloride and tetramethylamnioniuin triiodidc. JOL'KNAL, 73, 88 (1951); (d) J. Kleinberg, E . Colton, J . Sattiznhn and The teirabutylammonium tribromide and the tetraprop).lC. 4. VanderWerf, ibid. 76, 447 (1953); (e) C . Reed a n d R. S hlullianimoniuni iododibromide were, respectively, furnished 11y ken, i b i d . , 76, 3869 11954). L . Harris and S.E. Skelly of this Laborator!-. The purity of preparations \vas checked by iodometric titrations anti b>DEPARTMENT OF CHEMISTRY comparison of the melting points with those in the litcraSTATE UNIVERSITY OF IOWA ture,? ailti satisfactory agreements were obtained i t 1 each IOWA CITY, IOWA case. Considerable difliculties were experienced in t h e prcparation of the tetramethylammonium diiodochlcride which, t o the authors' knowledge, has not been, as yet, described Studies on the Chemistry of Halogens and of in the literature. The addition of one mole of iodine t o Polyhalides. V. Spectrophotometric Study of Poly- tetramethylammonium chloride in methanol solution rehalogen Complexes in Acetonitrile and in Ethylene sulted in the formation of a misture of the ciiiodochloride, tetraiodochloride and of pentaiodide ions. Somewhat better Dichloride results were obtained in using isopropyl alcohol a s a solvent. Upon cooling the solution containing a mixture of iodine a n t l BY ALEXASDER I. P O P O V A S D ROBERTI;. SI4 E N S E N tetramethylammonium chloride, first a mixture of the tetraRECEIVED MARCH 2, 1955 iodochloride and of diiodochloride was obtained, but up011 and letting the solution stand overnight, L: crol, of Absorption spectra of several polyhalogen coni- filtering bronze-red crystals was obtained which anal>-zcd t; Ix the plexes have been recently determined in various diiodochloride. The crystals had a m.p. of 102 . The solvents,' but only in the case of the iododichloride compound ib rather unstable and slowly gives up iodiiii.. and iodotetrachloride ions in acetonitrile, has the Complete decomposition was observed after several weeks. 'The absorption spectra of the diiodochloride solutions were investigation been carried out with thoroughness, determined on freshly prepared samples. and the absorption characteristics have been deterSolutions.-Preliminary experiments have int1ic:itetl that mined. This work was undertaken in order t o ob- the solutions of polyhalogen complexes in purified x c t o tain precise data on the absorption spectra of the nitrile and ethylene dichloride are quite stable for at least 24 hours. The solutions werc prepared by dissolving a ueighcd more common of the remaining polyhalide anions anlount of a salt in ;t given volume of the solvent anti thcti such as the IBr2-, IeBr-, IBrCl-, IzCl-, 13- and tlilutirig t o required concentrations. *411measurcmcnt w ( n Br3-. I t was of interest t o determine the absorp- made oil solutions prepared on the s;tnie day. Suppression of Dissociation.--Since it i h kno\vn t h a t t h c tion spectra in polar and non-polar solvents. For ions dissociate in acetonitrile, alld siuce it W A \ the first type, acetonitrile seems t o be a good choice polyhalide desired t o obtain the true molar absorbanc!- indices for because of high dielectric constant (36.5 a t 25"), these species, the absorption spectra were d fair resistance to halogenation, and especially its solutions contaiuiiig i t t i e.;cess o F t h e disso transparency down t o '300 nip. The choice of the ion i i i the form of the corriyxinditlg tetrrinicLtl 'rile hali,le wlt \vas atlrletl until t t o furlhei- chLitigci i i second solvent w m primarily determined by the ,31t. t11r :Ibsorptioii spectra W L I ~ubser\ e(1. S o i i c . IJT t lie hiiiiplc limited solubilities of ionic polyhalides in non-polar I~alitlrsa b s x b e d ~ p p I r c i d A ) .it tltc \\ .ive lcupths i~~!.e\tiliquids. The best possibility seems to be ethylene gated. dichloride (D = 10.2 a t 25") which can dissolve sufResults and Discussion
ficient amounts of the tetraalkylamrnonium polyhalides for spectrophotometric measurements and is trailsparerit dowri t o 2;Lj i n p .
(1'8 (a) A. I Popov a n d Ii €1. Schmorr, T H I S J O U K h . 4 1 , . 7 4 , .&Iii? , 1 9 3 2 ) ; fh) A I . Popov axid J X j e s s u p , i b i d . , 74, (!I27 (1!)52) c \ I
