1340
ALEXANDER I. POPOV AND DAVID H. GESKE [ C O S T R I B U T I O X FROM THE
DEPARTMENT O F CHEMISTRY, STATE UNIVERSITY O F
Studies on the Chemistry of Halogen and of Polyhalides. Iodine Species in Acetonitrile BY
Vol. 80
.kI,EX.INDER
I. P O P O V
AND
XIII.
IOWA]
Voltammetry of
DAVIDH. GESKE’
RECEIVED .4rGUST 12, 1957 T.oltammetric studies of iodine, sodium iodide and tetramethylammonium triiodide were made with acetonitrile as solvent Two of the waves obserx-ed for these species are well-characterized as involving iodide iori-triiodide ion and triiodide ion-iodine electrode reactions. The two waves occurring a t more positive potentials probably involve oxidation of iodine t o cationic iodine species. From the voltammetric data for equimolar mixtures of triiodide ion and iodide ion, and iodine and triiodide ion, the formal electrode potentials for the respective half-reactions were calculated. A study was made of the oxidation of iodide ion in the presence of pyridine, 2,6-lutidine, Zpicolirie, 2,2’hipyridine, 4,4’-bipyridine, 2,2’,2”-tripyridine and 1,lO-phenanthrolihe. The voltammetric reduction of dipyridine iodine (1) perchlorate also was studied. a n d a rotating platinum indicator electrode.
Introduction The aqueous iodine system has been the subject of a number of voltammetric studies.?-j The physical and chemical properties of acetonitrile make it an excellent solvent for a voltammetric study of the iodine species. Recently Kolthoff and Coetzee6made a preliminary study of the oxidation of chloride, bromide and iodide ions a t a rotating platinum electrode in acetonitrile. In the present work a comprehensive study of the iodine systern in acetonitrile was undertaken using a rotating platinum indicator electrode. Experimental Part Apparatus.-The same polarograph, polarographic cell and saturated silver chloride reference electrode were used as previously described.’ Once again, all potential measurements refer to the acetonitrile silver, silver nitrate (0.01 M ) p r i m a r y rpference electrode. The indicator electrode was a platinum bead, with a n area of 0.29 mm.2, fused into a soft glass tube. When rotated at 594 r.p.m., the linear peripheral velocity was 15.6 cni ./sec. The platinum electrode was cleaned by immersion in hot concentrated nitric acid, repeated rinsing with distilled water, immersion in acidic ferrous ammonium sulfate solution for several minutes and further rinsing with distilled This procedure gave a “clean” electrode which was subjected t o prepolarization as indicated in later sections, usually preanodization in the test solution at 12.4v. for one minute. Limiting currents observed with the rotating electrode are conveniently expressed in terms of the limiting current where 11 = i l ( o b s . ) / C m ~ ,and has the dimensions, constant 11, microampere liter/millimole. The 1: value is an arbitrary constant for a particular electrode reaction and is dependent upon the area of the electrode and the hydrodynamic conditions existing a t the electrode. Since the same electrode and rate of rotation were used throughout, comparison of various electrode reactions was made by means of the limiting current constant 11. i\ number of runs also were made with a gold-plated platinum bead electrode. An adherent deposit of gold was obtained by electrodeposition at 70’ from a solution of auric cliloride containing potassium cyanide and sodium dihydrogen phosphate. (1: (a) Abstracted in part from t h e P h . D . Thesis of David H. Geske, State University of Iowa, 1857. (b) A-ntional Science Foundation Predoctoral Fellow, 1055-1957. ( 2 ) n’. Kernst and E. J. Merriam, 2. p h y s i k . Che7n.. 53, 235 (1905). (3) E. Brunner, ibid., 56, 321 (1906). (4) According t o Bibliography of Publications Dealing with t h e Polarographic Method, J. Heyrovsky, Coilectiu:? of Ceech. Chem. C o m m s . , 2 1 , Supplement 1, 5 (1956). t h e oxidation of iodide, bromide and chloride ions was studied by E . 51. Skobets, N. N. Atamanenko :in I N E O N \ ~ o L l A h . r M E T R Y OF T E T R A M E T H Y L A M M O N I U M T R I I O D I D E , n'ave
I
Ppridirir
El/%,
i.
Em,
concn.
V.
Fa.
V.
0.0 .002 .004 .02
-0. l(i - ,I9 - .17 - .If - . 14 - .l2 -. .16 - .14 - .20
1.15 1 .OS 1.19 1.11 1.10 0.96 1.03 1.01 0.78
0.30 . 30 .30 .30 , 20 .2,5 .20 .17 .09
.or)
.1
.5 1.0 3.0
Wave I1
I + wave
-i,
e". 0.63
,544 .56$) . Glg . G3 .646 .676 . 70 .73
t o obtain the data given iii Table 111. A satisfactory forward trace could be obtained only if the trace was started a t a potential no more positive than the potential corresponding t o the beginning of the limiting current plateau for the I + wave. The suggestion of formation of an electrode film was verified by attempting a reverse retrace of the current-voltage curve immediately after the first trace. The current was greatly lowered throughout. Only when the electrode was cleaned in hot i i i tric acid could tlic original curve bc reproduced. A 2 niild pyridine solution showed the same resid-
V.
i).74
.71 , 03 . G1
-i, pa.
0.213 f338 .42fj .595
. 63
.625
.53 .5l .46
.661 .70 .7G5
1.3
it -
it I
1.82 1.9.: 1.93 1,80 I . 75 1.40 1.52 1.43 1.06
ual current as the supporting electrolyte, so no gross oxidation of pyridine could have taken place. Tllc exact composition of the interfering film is not known, but it probably is the product of the destructive oxidation of pyridine, perhaps involving the [ I ( P Y ) ~ ] +species. Rogers and Lordz3 encountered a similar situation in the oxidation of phenylenediamines on platinum and gold electrodes. They state t h a t insoluble oxidation and/or polymerization products apparently formed on the electrode. ,411 electrode would give a well-defiiicd (23) L U. Rogers and S. S . Lord, Jr., A n d . Chcm.. 26, 281 (1954).
March 20, 1958 1
0
PYRIDINE
0.5
A
1347
VOLTABIMETRY O F IODINE SPECIES IN ACETONITRILE
2,s-LUTIOINE
0 4,4-B1PYRIDINE
OS5
2,Z~IPYRIDINE
t
i
I
I
/I
I
0.1
‘
0 I
il0
I
0.05
I I I
A
E
+
4
u
I
0.01
‘ j_ 8
0.005
1
O*Oo6
I
/
/
d
/
V
I
/
0’
0.001
0
0.2
0.4
0.6
0.8
1.0
i +/iI,II. Fig. 7.-Amine
molarity (logarithmic scale) us. current ratio,
wave the first time i t was used, but subsequent polarograms became successively poorer. The current decrease caused by this film formation prevents making an accurate statement on the behavior of wave I V upon addition of pyridine. The presence of pyridine also influences the voltammetric behavior of triiodide ion. Data are given in Table I V and in Fig. 8. It was of interest to investigate the voltammetric behavior of iodine (particularly the oxidation behavior) in the presence of pyridine. The problem is complicated, however, by the fact that acetonitrile solutions of iodine to which pyridine is added, rapidly form appreciable amounts of triiodide ion. The amount of triiodide ion increases with time. Thus in a freshly prepared 0.137 m M iodine solution, 0.05 M in pyridine, approximately 27% of the total amount of iodine was present as triiodide ion; 24 hr. later, the amount of triiodide ion was 50y0of total iodine. Despite this inherent complexity of the sample solution, a voltammetric study was made. The current-voltage curve of a freshly prepared 0.137 m M iodine solution 0.05 M in pyridine showed only a small anodic current in the potential region 0.4 to I .5 v. Voltammetry of Dipyridine Iodine (I) Petchlorate.-The reduction behavior of dipyridine iodine (I) perchlorate is shown in Fig 9. The data in Table I are the average values for four determinations. The current increase labeled “C” cannot properly be described as a wave. The “peak” appearance is caused by the current decrease previously described. The current “constant” for the peak shows considerable variation,
3.6
-
3.2 2.8
-
2.4 2.0 1.6 -
1
1.2 -
0i
L
-
-
-
-
-
L
--LI-LI_-I-
1
- 0 . 6 -1.0 - 1 . 4 -1.8 E US. AgNOs, V. Fig. 9.-Voltammetry of dipyridine iodine ( I ) perchlorate, 0.416 miM, 0.1 M lithium perchlorate. Current, fia. vs. potential of RPE, volts vs. Ag, AgNOa electrode. Forward trace, anodized electrode. 0.6
0.2
-0.2
limiting ionic conductance of 104.5 for the perchlorate ion in acetonitrile. Thus the limiting ionic conductance of the dipyridine iodine (I) ion is approximately 75. Another positive iodine compound, diphenyliodonium perchlorate, showed no reduction waves a t a rotating platinum electrode. (24)
P. Wdden and E. J. Birr, 2. p h y s i k . Chem., 144, 269 (1929).
ALEXANDER I. POPOV AND DAVID H. GESKE
1348
Effect of Chloride Ion on the Oxidation of the Iodide Ion.-Kolthoff and Jordan6 found that iodide ion in aqueous solution 0.1 M in hydrochloric acid and 1 M in potassium chloride was oxidized to iodine monochloride. Figure 10 shows in curve 2 the effect of chloride ion on a 0.098 m M tetraethylammonium iodide solution. The arrows mark the height waves I and I1 would have had in the absence of chloride ion. It is seen that there is considerable current beyond that due to the oxidation of iodide ion t o iodine. It is presumed that the reaction
+ c1- - c -
1/J*
-4 IC1
occurs t o some extent.
oxidation of iodide ion to iodine in acetonitrile involves a two-step wave. This difference in voltammetric behavior would seem to be quite largely due to the greater stability of triiodide ion in acetonitrile as compared with water. Katzin and GebertZ5found a dissociation constant for the dissociation o€ triiodide ion into iodide ion and iodine of 1.01 X in aqueous 0.1 M perchloric acid a t 27.2'. The at dissociation constant in acetonitrilez6is 1 X 25O. It was of interest to derive the equation of the polarographic wave for the electrode reactions corresponding to waves I and 11. Previous workers5 have pointed out that for the reduction reaction Ip
O
Vol. 80
+ 2e- -+
21-
the reversible wave is given by the equation E
=
const.
i? + 0.03 log ___ - i) (id
It may be shown that for the electro-reduction corresponding to wave I, (I3- 2e-+ 3 I-), that
+
E
=
R - 0.03 log --r
(il
i3
- 2)
while for the electrooxidation E =R
2.s
iL ' 0.6
4
I {
/
I
1.2
1.4
J
0.4
0.2
E . V S . AgNOa,
0.0
V.
Fig. 10.-Effect of chloride ion on iodide ion oxidation. (1) Residual current, 0.1 M tetrabutylammonium chloride (Bu,SCl), left abscissa; (2) 0.098 m M tetraethylammonium iodide ( E t 4 S I ) ,0.01 M BuaSC1, 0.1 111 lithium perchlorate, right abscissa; (3) 0 218 m M EtaNI, 0.1 M BudSC1, left abscissa. Current, pa. us. potential of R P E , volts us. Ag, &l\To3 electrode; reverse trace.
m'here 0.1 I?!! tetrabutylammonium chloride was used as the supporting electrolyte, curve 3 was obtained. There is no well-defined wave, only an inflection point. The current a t the inflection point is approximately twice that which would be expected for the oxidation of iodide ion to iodine in perchlorate ion supporting electrolyte.
Discussion Waves I and 11, Iodine System.--The electrode reactions which occur a t wave I and I1 are unambiguously defined by the voltammetric data reported in the foregoing sections. They are wave I GI- - 4e- --+ 218wave I1 213- - 2e- +31%
The result of the millicoulometric study, n = 0.95 faraday/mole, is in satisfactory agreement with the above statements, ie., a one-electron change for the over-all reaction on a molar basis. It is interesting to note that Kolthoff and Jordan6 found only a single wave for the over-all oxidation of iodide ion to iodine in aqueous 0.1 Mperchloric acid solutions; as already has been shown,
+ 0.03 log (i, -i --
4 3
Evaluation of the half-wave potential from tliesc equations gives El/? = K'
- 0 06 log c
(5)
where C is the concentration of the electroactive material. It is expected that this wave equation would be applicable to wave I in the iodide ion oxidation and in the triiodide ion reduction. h plot of E vs. log !i/(il- i ) sfor wave I in the oxidation of iodide ion gave a straight line with a slope of 35 mv. compared with the slope of 30 mv. predicted from equation 3. A plot of E vs. log i3/(il- i) for wave I of the triiodide ion reduction gave a slope of -50 mv. instead of the predicted - 30 mv. The variation of half-wave potential of wave I of the iodide ion oxidation is shown in plate A of Fig. 11. Though there is considerable scatter from a linear relationship, i t can be seen that there is a t least a tendency toward more negative potentials a t higher concentrations of iodide ion. The solid line has a slope of - 30 mv. A similar plot is given in plate B for wave I of the triiodide ion reduction. The slope of the solid line is -70 niv. It is seen that equation .i predicts a slope of --ti0 mv. for a plot of El/, V S . log C. Using a procedure analogous to that used to dcrive equation 4, the equation for a wave corresporlding to wave 11, (213- - 2 e - 4 312)is obtained
while for the corresponding reduction
Evaluation of the half-wave potential gives El;?
=
k'
+ 0.03 log C
(8)
( 2 5 ) 1. I. Katrin aud E. Gebert, THISJOURNAL, 76, 204!3 (1054) (26) N. E. Skelly, Ph.D. Thesis, S t a l e University of Iowa. 195.5
March 20, 1958
VOLTAMMETRY OF IODINE SPECIES I N ACETONITRILE
1349
A plot of E us. log (-i)3/(i1- i)zfor wave I1 of basis for purposes of comparison. From previous the triiodide current-voltage curve gave a straight conductivity work, 22 the limiting equivalent conline with a slope of 31 mv.; a plot of E vs. log ?/ ductances of tetramethylammonium iodide and (il - i ) 3for wave I1 in the iodine reduction gave a triiodide are 195.7 and 191.5. Taking the limiting straight line with a slope of - 34 mv. ionic conductance for the tetramethylammonium Plates C and D in Fig. 11 show the vari-0.04 -0.020.00.020.04 0.32 0.30 0.28 0.26 0.24 ation of half-wave potential with depolarizer concentration. Data in plate C for wave I1 of the triiodide ion oxidation a t 10-3 least show a tendency for a shift of halfwave potential to more positive values a t 10-3 higher concentrations. The solid line indiI cated has a slope of 44 mv. A similar plot I r“ v for wave I1 of the iodine current-voltage v curve (plate D) more closely approximates a linear relationship. The solid line has a 10-4 slope of 55 mv. Discussion of the results of the voltain10-4 metric studies of the iodine system in terms of equations 3-8 does give a basis for the ! O concentration dependence of the half-wave 10-3 potentials. Because of the reality of the influence of electrode surface condition, it does not seem that a completely valid evaluation of the “reversibility” of the electrode H” v reaction can be made from the logarithmic analyses of the waves, which show some 10-4 deviation from the predicted slope values. I t is seen from the voltammetric data for iodide-triiodide ion and iodine-triiodide mixtures (Fig. 5) that there is no iuj2ection 0.22 0.18 0.14 0.10 0.06 0.32 0.30 0.28 0.26 a t the zero current point for either wave I E 2’s. :Ign‘Oa, V. E V S . -4gso3, V. or wave 11. On the basis of this observation, the conclusion is reached that elecFig 11.-Variation of half-wave potential with concentration. trode reactions corresponding to waves I Jt’ave I : plate A4,iodide ion oxidation; plate B, triiodide ion and I1 are “reversible.” The apparent reduction. XVave 11: plate C, triiodide ion oxidation; plate D, deviations from reversible behavior are iodine reduction. Concentration of depolarizer (logarithmic scale) considered to be due to electrode surface 21s. half-wave potential, volts, us. -4g, AgN03 electrode. effects. From the same data (Fig. 5) it is possible to esti- ion in acetonitrile as 92.9,25the limiting ionic conmate the formal electrode potential for these half- ductances of iodide and triiodide ions are 102.8 and reactions. Since both oxidized and reduced forms 98.6, respectively. It is seen that the sum of the are present in solution in equal concentration, the equivalent current constants is essentially the same zero current potential can be related to the formal for iodide ion as for triiodide ion, as would be expected on the basis of their similar diffusion coefelectrode potential. ficients. For the half-reaction The diffusion coefficient of iodine in acetonitrile 1 3 - + 2c - --+ 31 is not known. Consequently no effective comment (CI ) 3 can be made on the fact that the equivalent current b = b” 0 03 log (CY8 constant for iodine is larger than the constant for For the equimolar mixtures, CI- - CI~.. = 4.5X both iodide and the triiodide ion. -11, E , = 0 = -0.07 v., thus E” = -0.27 V . W . The decrease of the ratio of the current constants Ag, AgN03 electrode. -In analogous calculation of waves I and I1 in the series iodide ion, triiodide for the half-reaction ion, iodine is indeed puzzling. From examination of the individual equivalent current constants it 312 2e- --+ 213appears that it is the increase of the current congives a value of 0.39 v. StrehlowZ7calculated the stant for wave I1 (in the same series) that is restandard electrode potential for the iodine-iodide sponsible for the decrease of the ratio, particularly half-cell in acetonitrile from solubility measure- for iodine. No reason can be advanced to explain ments of sodium iodide. The reduction potential this situation. was calculated to be -0.16 v. vs. Ag, AgN03 elecWaves I11 and IV. Iodine System.-Waves I11 trode, a value intermediate to the formal potentials and IV observed in the voltammetric studies of the (in 0.1 11.1 lithium perchlorate supporting electro- iodine system must represent oxidation of iodine to lyte) given above. some positive valence state. It seems reasonable to Current constants are listed in Table I on both a disregard the oxygen-containing positive iodide ions molar and normal (termed “equivalent” constants) as possible reaction products because of the non(27) H. btrehlorv, d L l e i f i d i c > i i 56, S 2 7 (19.52) oxygenated nature of the solvent and because it was h
h
Y
h
+
1350
ALEXANDER I. I’OPOV
found that addition of 0.5y0water had no effect on the oxidation of iodide ion. A number of positive iodine compounds of valences one and three are known.%2%29 The current constant for wave I11 is grossly smaller than would be expected for the usual oneelectron reaction. I t is concluded that abnormal effects must be in force. The fact that the temperature coefficient for wave I11 is not grossly different from the temperature coefficient for waves I and I1 (which are assumed to be diffusion-controlled) strongly suggests that wave I11 is not a “kinetic” wave. I t is seen in Table I1 that the millicoulometric data do not clearly distinguish a difference between the sum of reaction for waves I and I1 and the sum of waves I, I1 and 111. The IZ values, 0.95 and 0.98 faradays/mole, respectively, coincide within experimental error with the expected value of 1.00 for the sum of waves I and 11. Perhaps the most significant experimental obscrvation with regard to wave I11 is the fact that it does not occur in the current-voltage curve of molecular iodine, but i t is found in the oxidation of triiodide and iodide iocs. These observations suggest that wave I11 occurs only when nascent iodine is produced a t the electrode surface. It appears from the study of the effect of pyridine on iodine species that nascent iodine is also necessary for oxidation of iodine to the pyridine-stabilized monovalent cation (vide infra). The similarity of this case with the behavior of wave I11 is further intensified when it is noted that very low concentrations of pyridine, 5 X 31, apparently shift wave I11 in the iodide ion oxidation to more negative potentials (indicating a stabilization of the reaction product). However, the conclusion cannot be drawn from these data that the two waves necessarily represent the same electrode reaction.
It i s not possible to deteriiziiae urzambiguously the natuye of the electrode reaction at wave IIT; the best postulate seems to be that oxidation proceeds to the monovalent iodine cation by a mechanism that requires nascent iodine. Wave IV occurs as an oxidation wave in the current-voltage curves of iodine, iodide ion and triiodicle ion in acetonitrile. The fact that the current constants are dependent upon concentration (Fig. 4) suggests that the over-all electrode reaction is of a complex nature, perhaps involving more than one primary reaction; apparently a t low concentration one of the primary electrode reactions makes a larger contribution to the total current than would be expected from a linear current-concentration relationship. The possibility exists that the electrode cannot be considered totally inert. If electrode dissolution or interaction contributes to the total observed current a t wave IV i t is a t least possible that the over-all current constant would not be independent of concentration, i.e., the limiting current would not necessarily be strictly proportional to concentration. (2s) Gmelin’s “ R a n d h u c h der Anorganischen Cheniie,” St!] Ed., Verlag Ctremie. Eerlin, 1933, System No. 8, p. 454 if. (29) J . Kleinberg. “Unfamiliar Oxidation States and Tlreir Stahilira. Lion,” University of Kansas Press, I.awrence, Kausas, 15150, p . 37 fl.
AND
DAVIDH. GESKE
VOl.
so
The similarity of a-values, which characterize the rate of rotation dependence for waves I , I1 and IV, indicates that oxide formation does not make an appreciable contribution to the current a t wave IV. Under the criteria stated by previous workers30 the a-value for wave IV also indicates that wave IV does not involve kinetic currents. The magnitude of the current constant for wave IV may be altered by an electrode “interaction.” The fact that a t higher iodide ion concentration the current constant approaches a limiting value of --G pa. liter/mmole, in comparison with a current constant of -4.91 pa. liter/mmole for the oxidation of iodide ion to iodine, gives some indication that the ultimate electrode reaction is oxidation of iodine to the monovalent state. ,‘I further consideration of the rnillicoulometric data for electrolysis of iodide ion a t 2.11 v. (Table 11, run C) is given in Table V. TABLE V ~ ~ l l . L I C O U L O M E ~ r R ISCT U D Y O F E L E C T R O D E REACTION, \\’AVE
Iv Electrode reaction
Electricity. faradays
0.327
Residual current Waves I, 11, I11 LVave I V
0.05 0.006“ 1.04
0.630
Total observed
1,696
I,-,
p
moles
.__
-
Estimated from height of limiting current for waves I, I1 and 111.
The detection of positive iodine species produced electrolytically is based on their eventual reaction with iodide ions to form triiodide ions. Thus in the above listing, from the quantity of electricity consumed by reaction IV, 1.04 pfaradays, one would expect to find 1.04 pnioles of triiodide ion if the product of the electrode reaction were the I + species and if all of the I + underwent the oxidation reaction with iodide ions. The experimentally observed amount of triiodide ion was 0.327 pniole, COIIsiderably less than the expected 1.04 pmoles. I t is possible that the product of the electrode reaction a t wave IV (presumably I+) not only reacts with iodide ion but also undergoes a side reaction, perhaps with the solvent. In the case of the reaction correspolldirlg to wave IV, it i s impossible to state unambiguously the exact nature of the electrode reactioia. The following seems to be a sound postulate: wave IV represents a complex electrode reaction involving the oxidation of iodine to a t least monovalent positive iodine. Accompanying electrode interaction possibly causes the non-linear relation between limiting current and concentration and the high value for the current constant for wave IV. Effect of Pyridine on Oxidation of Iodine Species.-The occurrence of a third wave in the current-voltage curve of iodide ion in the presence of pyridine is rather conclusive evidence for oxidation of iodine to the monovalent cation. The product of the electrode reaction is assumed to be dipyridine iodide (I), but the inotiopyridine iodine (I) ion is also k ~ i o w ~ ~ . ~ (30) S. S. Lord, Jr., a n d I,. 13. Ropels, A n a l . Chcm., 26, 284 ( I i J b i ) .
March 20, 1955
VOLTAMMETRY OF IODINE SPECIESIN ACETONITRILE
The plot of current ratio w. pyridine concentration for the iodide ion oxidation, as given in Fig. 7, approaches a value of one as would be expected for equivalent one-electron reactions as I- - e - + 1/*1* l/tL
+ ~ P Y- e -
-
[I(PY)z~ +
The presence of pyridine, even up t o 3 J4 concentration, had substantially no effect on the magnitude of waves I and 11. The half-wave potential of wave I1 was shifted t o more negative potentials with increasing pyridine concentrations. This fact apparently is responsible for the merging of waves I and I1 a t high pyridine concentrations. The derivation of a relationship between pyridine concentration and the half-wave potential of the positive iodine wave is rather difficult since it involves treatment of a consecutive wave. It is sufficient to note that a shift of the half-wave potential for the positive iodine wave (for both iodide ion oxidation, Table 111, and triiodide ion oxidation, Table IV) t o more negative potentials with increasing pyridine concentration was observed. The effect of pyridine on the oxidation of triiodide ion was generally analogous to the electrooxidation of iodide ion in the presence of pyridine. However, it is seen in Fig. 8 that the plot of current ratio, i+/i1,11vs. pyridine concentration, does not aproach a value of one, as was the case for iodide ion. The limiting value ca. 0.4 is grossly less than the value of 0.95 found with the iodide ion. This indicates that relatively less iodine is being oxidized to the monovalent state with triiodide ion than with iodide ion. The conclusion drawn from the plot of current ratio, i + / i I r , vs. pyridine concentration (Fig. 8) where ~ I isI the oxidation current of triiodide ion, is that only as much iodine is oxidized t o the plus one state as is produced electrolytically by oxidation of triiodide ion, where the oxidation is considered to proceed through dissociation of triiodide ion into iodine and iodide ion. These observations suggest that nascent iodine is a prerequisite of oxidation to the monovalent state. A further test of this statement would be offered by investigation of the voltammetry of iodine in the presence of pyridine. As was indicated earlier, the complex nature of acetonitrile solutions of iodine and pyridine (increasing triiodide ion concentration) prevents a totally rigorous voltammetric study. However, the absence of a n y gross oxidation current in the iodine-pyridine current-voltage curve in the potential region 0.4 to 1.5 v. offers support for the contention that nascent iodine is necessary for the oxidation of iodine to the monovalent cation in the presence of pyridine. A possible reaction scheme might be I-
- e--+
I*
,
+ 2py - e--+- k
$I*
I(py)~+
(9)
I2
where I * is atomic iodine a t the electrode surface. This scheme would qualitatively account for the dependence of the magnitude of the positive iodine wave on pyridine concentration, Le., at high pyridine concentrations, the probability for the oxida-
1351
tion of iodine “atom” may be increased in comparison with its probability for recombination with another iodine atom to give an iodine molecule (which is inert to further oxidation). The postulation of atomic iodine on the electrode is based on the work of Vetter3‘ who found that the reaction I +e-+
I-
was the potential-determining step in the establishment of the iodine-iodide oxidation-reduction potential on a platinum electrode. Effect of Other Heterocyclic Amines on Iodide Ion Oxidation.-It is seen from Fig. 7 that there is a distinct difference in the extent to which oxidation of iodide ion to the monovalent cation occurs in the presence of various amines. The fact that the current ratios decrease in the order pyridine > 2picoline >> 2,6-lutidiiie, strongly suggests that steric factors are of importance in the electrode reaction. This assertion is further supported by the observed current ratios for the polynuclear amines. Thus 4,4‘-bipyridine stabilizes the iodine cation almost as effectively as 2-picoline, whereas the formation of the iodine cation occurs to a very small extent in the presence of 2,2’-bipyridine. The 2,2’bipyridine molecule may presumably exist in solution in either the cis or trans configuration. Fielding and L e F & ~ rfound e ~ ~ by dipole measurements that the 2,2’-bipyridine molecule exists in the trans configuration in benzene solutions. In either configuration, cis or trans, 2,2‘-bipyridine would show more steric hindrance to reaction with iodine species than would the 4,4’-isomer. It should be acknowledged that the apparent stability of the 4,4’-bipyridine iodine (I) complex may be due in part t o a tendency toward p o l y m e r i ~ a t i o n . ~ ~ Brown, et ~ l . have , ~ ~ pointed out that base strength must be considered a function of the steric requirements of the acid. The order of base strength toward hydrogen ion is pyridine > 2-picoline > 2,E-lutidine. There is an inversion to the reverse order with an acid such as trimethylboron. Thus i t is difficult t o separate clearly basicity effects from steric considerations. If the assumption is made that the electrode reaction scheme is analogous to equation 9, i t seems qualitatively correct to suggest that the steric characteristics of the amine affect the rate constant, k, which in turn influences the observed current. Such an approach would agree with a low current ratio in the case of iodide ion oxidation in lutidine solutions. The pronounced decrease in observed current might then be attributed t o the “blocking” effect of the two methyl groups in a reaction between the amine and the iodine atom. This would be particularly true if it were required that two amines coordinate with the iodine. The true base strength might be expected to determine the stability of the iodine amine complex. This stability would in turn influence the half-wave (31) K.J, Vetter. 2.Elcklrochcm., 66, 121 (1951); 2. physik. Chcm., 199, 22, 285 (1952). (32) P. E. Fielding and R . 3. W. LeF$vte, J . Chcm. Soc.. 1811 (1951). (33) F. B. Stute, Ph.D. Thesis, State University of Iowa, 1957, p.
109.
(34) € C. I.Brawn, D. Gintis and L. Domaeh, TEISJOURNAL, 7 8 , 6387 (1956).
1352
B. H. ECKSTEIN AND E. R. VANARTSDALEN
potential of the positive iodine wave. Perusal of the half-wave potentials in Tables 111 and IV does not show any clear pattern. It is concluded that the half-wave potential is not grossly affected by the base strength of the amine with the series pyridine, 2-picoline, 2,B-lutidine. Rather the voltammetric behavior is influenced to a greater extent hy the configuration of the amine. Voltammetry of Dipyridine Iodine (I) Perchlorate.-It is difficult to interpret the observed voltammetric behavior of dipyridine iodine (I) perchlorate because the limiting currents are substantially smaller than would be expected. From the current constant for the over-all reaction, I- e- + ‘/2 1 2 , -4.91 pa. liter/mmole, one would estimate an over-all current constant of 9 to 10 pa. liter/mmolo for the reaction35 [I(PY)zI +
+ 2e-
-----f
I-
+ 2 PY
The current constant for wave A is only 1.22 pa. liter/mmole. This low current constant can only be partly accounted for by a decreased diffusion coefficient. Cases of similar abnormally small current constants were found in the interhalogen and (36) T h e dependence of limiting current observed a t a R P E on the diffusion coefficient ranges from 0.5 t o 1.0 (il = ! 2 D 0 ~ J - - f ~ 9From , the limiting ionic conductances of t h e dipyridine iodine (I) ion and t h e iodide ion, 75 and 102.8. respectively, i t is evident t h a t even a direct proportionality between limiting current aud diffusion coefficient gives a ratio of only 0.74 between the current constant for the iodide ion t o iodine oxidation and t h e reduction of clipyridine iodine (I! t o iodine.
[CONTRIBUTION FROM THE
VOl. 80
mixed polyhalide ion r e d ~ c t i 0 n . l ~No satisfactory explanation has been found. The following electrode reactions are tentatively proposed to explain the current-voltage curve observed for dipyridine iodine (I) perchlorate.
+ +
6 I ( p ~ ) ~ +Ge- --+ 312.pyz wave B 312 py2 3e- --+ 2132py wave C 2132py 4e- +612py wave A
+
+
+
+
I n view of the abnormally low current constant of wave A, the current constant for wave B, 0.7 pa. liter/mmole, seems qualitatively satisfactory for the reduction of iodine to triiodide ion (the observed current constant for the reduction of iodine to triiodide ion is 2.11 pa. liter/mmole). It has not been possible to establish definitely the nature of the electrode reactions taking place in the electrochemical reduction of dipyridirie iodine (I). The electrode reaction tentatively suggested for wave A probably involves an appreciable overvoltage because its half-wave potential is even slightly more negative than that for the reduction of iodine to triiodide ion. It would seem reasonable to suppose that it would require less energy to reduce an electronegative group like [I(py)*]+than to reduce molecular iodine. This irreversibility is also found in the electrode reactions postulated for wave B and peak C. IOIVACIry, IO\\A
DEPARTMENT OF
CHEMISTRY, CORNELL UNIVERSITY]
The Heat of Formation of Boric Oxidel BY BERNAKD H ECKSTEIN~ AND ERVINR. VANARTSDALEN~ RECEIVED AUGUST 28, 1957 T h e heat of formation of amorphous B ~ 0 is 3 304.6 kcal /mole as determined by burning crystalline boron in oxygen without the use of any organic initiator. This agrees well with the value of 302 kcal./mole found by Prosen using the decomposition of diborane as the experimental route.
Introduction This paper describes the determination of the heat of formation of largely amorphous boric oxide by the direct union of elementary crystalline boron and oxygen without the use of any organic initiator. The heat of formation is defined as the enthalpy change for this reaction 2B(cryst)
+ 3/2 0z(g)
=
BZOdamorph)
The earlier determinations of this heat of formation are summarized in Table I. Note the wide divergence in the results obtained whenever organic initiators were used. The agreement between the 304.6 kcal./mole found in this study, using the most direct method of determining heats of formation, and the very careful work by Prosen, et ai., (1) T h e work here reported was carried out a t t h e Department of Chemistry, Cornell University, during 1948 and 1949. It forms part of a P h . D . thesis submitted to Cornell in February. 1P.53, by B. €i. Eckstein. (2) Research Laboratories, National Carbon Company, Division of Union Carbide Corporation. P. 0 . Box 61 16, Cleveland 1: Ohio.
makes i t seem very likely that the heat of forination of B2O3 is indeed ca. 302 kcal./mole. Description of Apparatus The actual combustions were carried out in a double valve type, self-sealing Parr Oxygen Bomb of 330-rrli. capacity, made of Illium steel. This bomb was immersed in about 2600 g . of distilled water, weighed to 0.1 g., in a calorimeter consisting of three concentric cylindrical cam, all chrome plated and highly polished. There were an outer and an inner can and an interspersed aluminurn radiation guard. A11 three elements were fastened a t the top t o a “Lucite”’2 ring. The bomb in the inner can was completely submerged but did not touch bottom, so t h a t water could circulate freely about the bomb. The cover for the assembly consisted of a chrome-plated brass disc as the lower face, a moriel metal plate as the upper face and a one inch block of “Lucite” as insulating material in between. The cover was pierced t o permit access for the various leads t o the bomb, the resistance thermometer, anti a platinutn resistance thermometer for calibration purposes. The pump type stirrer consisted of a shaft with four four-bladed propellers. The stirrer shaft was made in two sections joined by a section of “Lucite” to minimize heat losses. The pu~iip housing served as support for a four-lead copper resistance thermometer. Because the calorimeter was completely submerged in a constant temperature bath a superstriict ure
