NOTES
Dec. 5 , 1955 The equilibrium constants for direct salt formation, reactions 3 and 4, were calculated from the intensity of the acetic acid monomer carbonyl band and the acetic acid dimerization ~ o n s t a n t Tables ,~ I and 11. As before, reaction 1 appears almost quantitative and its equilibrium constant, therefore, could not be satisfactorily calculated. Reaction 2 is also nearly quantitative but constants of 150 =t50 1. mole-’ were approximated for the reaction in both solvents. No attempt was made to calculate a dimerization constant since the solvent and the excess amine interfere with the bands necessary for the calculation. A comparison of the equilibrium constants for direct salt formation is shown in Table 111. Although the base strength of the amines is expected to decrease for the series studied, the constants in CCL indicate an apparent increase. This enhancement of apparent base strength probably reflects the increase in the ability of the alkyl ammonium ion to solvate the acetate ion with the formation of a diprotonic bond in diethylammonium acetate and a triprotonic bond in n-butylammonium acetate. In CHC13 this effect is superimposed on the ability of the CHCla to solvate the ion-pairs and the amines. The latter effect has been determined quantitatively as the heat of solvation of the TABLE I THE EQUILIBRIUM CONSTANT FOR REACTION 3 (ACETICACID0.0008 M ) c
M
CdHoNH2 10g(’O’l) 5.88~
0.0000 ,0008 ,0016 .0020 .0032 ,0040 .0048 ,0060 .0080
0.359 .288 ,237 ,212 ,176 ,128 .I30 .098 ,071
cx
HOAc
(HOAc):
(6.37) 5.11 4.20 3.77 3.12 2.28 2.31 1.74 1.26
(0.81) .52 .35 .28 .I9 .10 .ll .06 .03
IN
CC,
10AcOHaNC4Ho
C4H9NHz
K
..
...
..
1.85 3.10 3.67 4.50 5.52 5.47 6.14
6.15 12.90 16.33 27.50 34.48 42.53 53.86 6.68 73.32 Av.
590 570 600 530 700 560
660 720 600
TABLEI1 THE EQUILIBRIUM COXSTANT FOR REACTIO (ACETICACID 0.0008 M ) M log ( I o / I ) GH~NHI 5.71~
HOAc
c x 104 AcOHa(H0Ac)n NCdH9
0,0000 .0008 .0016 .0024 ,0032 .0064
(3.40) 2.44 1.73 1.21 0.91 0.45
(2.30) 1.19 0.60 .29 .I7 .04
0.220 .I56 .112 .078 .059 .029
C4H9NHa
K
..
...
..
3.18
4.82 10.93 17.79 25.25 56.53 Av.
2650 2680 2880 2940 2930 2800
6207
amine in CHC134, Table 111. Thus ion-pair structures and solvation play a prominent role in determining the apparent base strength of amines in non-dissociating solvents. Examination of the monomer ion-pair structures presented for systems containing acetic acid and primary, secondary and tertiary amines indicates that hydrogen bonding occurs between the alkylammonium ion and one oxygen atom of the acetate ion. For salts of primary and secondary amines, where more than one acidic hydrogen atom is available, multiple hydrogen bonding to a single oxygen atom occurs. I n the more concentrated solutions of the primary and secondary amine salts, the multiple hydrogen bonding of the ions appears to lead to the formation of bridged dimers. (4) M. Tamres, S. Searles, E. M. Leighly and D. W. Mohrman, THISJOURNAL, 1 6 , 3983 (1954).
DEPARTMENT OF CHEMISTRY NORTHWESTERN UNIVERSITY EVANSTON, ILLINOIS
Polarographic Studies on the Bipyridines BY ANTOINEB. ZAHLAN AND ROBERTH. LINNELL~ RECEIVEDJUNE 13, 1955
No polarographic studies appear to have been previously reported on the six isomeric bipyridines.2 Pyridinium ion produces a catalytic hydrogen wave a t the dropping mercury electrode; in phosphate buffers (PH 6.1 or 7.4) pyridine is reduced (Ell2 = -1.69 vs. S.C.E.) apparently by a six-electron p r o ~ e s s . ~Smith4reduced bipyridines with hydrogen and platinum oxide in acid medium and also with tin and hydrochloric acid; increasing d i s culty of reaction was reported as 4,4’-, 2,4‘-,6 2,2’-, 33‘-. The reduction of 4,4‘- and quaternary bases derived therefrom has been of interest in developing oxidation-reduction indicator^.^,^ Experimental Materials.-2,2’-
was purchased from Eastman Kodak Co. and 4,4‘- from the British Drug House. 4,4’-, 2,2’and 2,4’- were prepared by the sodium-pyridine reactions and pyrolysis of pyridineByielded 2,3’- and 2,4’-. p-Phenanthroline (prepared by Skraup synthesis) was oxidized’o to make 3,3’-. Melting points of the bipyridines and their picrates plus Amax. and hax. ultraviolet values agreed well with literature values.436J1 Analytical reagent grade chemicals were used in preparing buffers which were all 0.10 M in
(1) Titanium Zirconium Co., Inc.,Flemington, N . J. (2) Z,Z’-Bipyridine, 3,3’-bipyridine, 4,4’-bipyridine, 2,3’-bipyridine, 5.07 2,4’-bipyridine, 3,4’-bipyridine. No work is reported here on 3,4’6.21 as we did not have a sample. For brevity we omit bipyridine in all subsequent references in this paper. 6.75 (3) P. C. Tompkins and C. L. A. Schmidt, J . Biol. Chcm., 145, 843 7.47 (1942); M. Shikata and I. Tachi, Mcm. Cull. Agr. K y o f o I m g . Uniu., 4, 19 (1927); E. Knobloch, Collection Czech. Chem. Communs., 12, 407 (1947); M. K. Shchennikora and I . A. Korshunov, J . Phys. Chem.. TABLE I11 (U.S.S.R.),2 2 , 503 (1949). (4) C. A. Smith, THISJOURNAL, 60, 1936 (1928); IS, 277 (1931). EQUILIBRIUM CONSTANTS FOR THE REACTION OF AMINES (6) Smith called this 3,4’- but Krumholz [THIS JOURNAL, 7S, 4449 WITH 0.001 M ACETICACIDIN CCL AND CHCls SOLUTIONS (1961)) has shown it is 2,4’-. AH4 (8) L. Michaelis and E. Hill, J . Gen. Physiol., 16, 859 (1933); Amine K(CC1S K(CHClr) (cal. mole-’) THISJOURNAL, 66, 1481 (1933). Triethylamine’ 800 3000 870 (7) H . Kuhn, Helo. Chim. Acta. SP, 2247 (1949). Diethylamine2 2800 3000 886 (8) C.A. Smith, THISJOURNAL, 46, 414 (1924). C. A . , 44, 399% (9) C. Krumholz. Selccfa Chim., No. 8, 3 (1949); 6. n -Butylamine 2800 600 714 (1950). Pyridineo 200 60 484 (10) C. A. Smitb, THISJOURNAL, 68, 397 (1930). To be published. (11) C. Krumholz, ibid., 1 3 , 3487 (1951).
NOTES
6208
VOl. 77
KCI. Buffers used: PH 2.7, Britton and Robinson'*; pH 4.0, 0.12 M potassium acid phthalate; pH 4.2, 0.40 ill HOAc and 0.25 A I NaOAc; p H 4.6, 0.20 M HOAc plus 0.20 AI NaOAc; pH 5.8,0.16 M NaHzPOa; pH 6.3, Britton and Robinson; pH 8.5, 0.16 ?\Ta?HP04; pH 13, 0.3 AI NaaPO& Equipment.-A Beckman DU was used for ultraviolet spectra. PH were measured on a portable Cambridge Instrument meter. Polarograms were photographically recorded on a Sargent-Heyrovsky Model XI1 instrument with a 10-ml. Heyrovsky erlenmeyer cell with mercury pool anode. The electrode characteristics were: t = 3.80 seconds/drop and m = 1.62 mg./sec., determined in distilled water, open circuit, 66.5 cm. mercury, 25'.
presence of 2,2'- acts as a maxima suppressor in mixtures of bipyridines but complexing with F e + + destroys the suppressor function. Extensive concentration-current information was not obtained for 2,4'- and 2,3'-. For solutions 0.22.0 X loF3 mole/l. (pH 4-5). The first wave is clearly defined and pre-waves appear to be absent. The Ilkovic equation is valid using the same constant as for 4,4'-. In neither case is the second wave sufficiently well defined to permit objective measurement. $3'- is difficult to reduce and its polarographic wave appears very close to that of Results the hydrogen discharge wave. Table I summarizes half-wave potentials. The 4,4'- shows the characteristic blue color of its redifficulty of reduction of bipyridine increases in the duced form a t the surface of the mercury drop when E1/2(I)13 is reached. Logarithmic wave plots for order 4,4'-, 2,4'-, 2,3'-, 2,2'- and 3,3'-. For anathe first wave give straight lines with slopes of lytical work on mixtures of bipyridines, acetate buf0.060. Therefore the polarographic work confirms fers pH 4-5 are suggested. 2,2'- should be cornearlier potentiometric6 and theoretical' work which plexed with FeS04 and determined spectrophotoindicated that a one-electron reversible reduction metrically. 4,4'- is determined from the height converts 4,4'- to the unstable blue product. The of its EliZ(Ijwave which is easily separated from calculated diffusion coefficient (first wave, assum- all the other bipyridine waves. I n a mixtureEl/,(I) of 4,4'would ing n = 1 and using experimental &/c value with of 2,4'-, 2,2'- (and 3,4'-14j and E1/~(11) Ilkovic equation) is 0.64 X 10-5 cm.l set.-' a t pH all superimpose ; since 4,4'- is determined separately 4.6 and 25'; the Stokes-Einstein value of Do is it can be subtracted yielding a value for 2,4'- and 0.71 X set.-' (using p = 1.21 g . / ~ m . ~ ,2,3'- (+3,4'-14) combined. Determination of 3,3'our experimental value). The second wave does is not satisfactory in aqueous solutions. not yield a linear log plot. The Ilkovic equation is T ~ R LI E obeyed for both waves, between 0.2-10 X T*ARIATION OF HALF-WAVE POTEVTIALS OF THE BIPYRIDINES mole/l. (above 2 X IOp3 mole/l. maxima must be n'ITH p1-I suppressed: 2,2'- does this) where &(I) = 2.67 X C X lop6 ampere and id (I 11) = 5.45 C X T = 25 & 2 " , concn. is of the order of 10-3 mole/liter unless otherwise indicated ampere (C = concn. in mole/l. x IO3) in ace-El/*(I) -EI/z(II) tate buffers (pH 4.62) at 25". The second wave pH 2 ?'-a 23I-b 2 4'-b 3 3'4,4'2,2'4,4'. seems also to involve one electron but further work 4 . 0 . . . . . . . . 0.88 .. .. is necessary to explain the non-Iinear log plot. 4 . 2 1.12 .. .. .89 1.24 . . From 16-31' the temperature coefficient of the first 4 6 1 . 1 4 1 0 0 1 04 1 .XC ,912 1 . 2 5 1 . 0 6 wave is linear and corresponds to 2.5% increase in 5 8 1.21 1.10 ,. .. .. .. ;,(I) per deg. C. 6.4 . . 1.14 . 1.06 .. 1.10 2,2'- has a small prewave whose height (1.14 X 8.3 ,. . . .. 1.31 .. .. 10-6 ampere) is independent of concentration, 8.8 1.46 . 1.20 .. .. probably due to strongly adsorbed reduction prod8 .9 .. . . . . 124 . 1 ..?4 uct. The prewave merges into a wave which is 9.1 1 .:if3 , . followed by a second wave; neither i(prewave .. 202d .. 5.0 .. . . .. I 11) nor i(I 11) is linear with concentration .. 1 30e 13.0 .. . . 1 38' .. although smooth concentration-current curves a Since a pre-wave merges into I, these E I / ~ ( values I) were may be plotted. The current increases too slowly with concentration; a 1.0 X inole,'I. solution determined where i (including pre-wave) = i total/2; they ) Although second therefore not true E > / ? ( I values. (acetate buffer) of 2,2'- has a current (after sub- are polarographic waves do occur, the definition of the waves is tracting the prewave) 30% greater than a corre- so poor that good El/*(II)values could not be estimated sponding solution of 4,4'-. This effect decreases for 2,3'- and 2,4'-. Concn. of 3,3'- was 10.82 X mole/ liter. Concn. of 3,3'- was 21.65 X 10-3 mole/liter. with increasing concentration until a t 3.0 X polarographic waves of 3,3'- owing to their occurrence mole,/l. the 2,2'- current is the same as that for aThe t very negative voltages are obscured by the hydrogen wave 4,4'-. E(1) shifts to more negative values above except a t higher concentrations. The value given for Ell2about 2.0 X mole//l. and a t 5.0 X mole/l. ( I ) appears to correspond to two electrons and is therefore E(1) has merged into E(I1) with the prewave no the combined waves I plus 11. e A t this high pH E I / ~ ( I ) longer visible. The prewave makes i t very difficult and E1j2(1I)are the same. to separate the first wave for a log plot. The secAcknowledgment.-The authors greatly appreciond wave yields a linear log plot with an apparent ate support of The Research Corporation of New n value of one. -4ddition of sufficient FeS04 to York City which made this work possible. form the complex, Fe(2,2'-)8++, completely reDEPARTMEXT moves the polarographic waves of the 2,2'-. The CHEMISTRY
+
-
-
, .
, .
,
+
+
,
+
(12) 0. H. Muller, "The Polarographic h l e t h o d of Analysis," C h e m . Ed. Pub. Co., E a s t o n , Pa., 1951, p . 104. (13) T h e symbol (I) refers t o t h e first polarographic wave, (11) t h e aecond, a s t h e merciiry elcctrt,de potential is mad? in,,re neg.itr\e,
AMERICAN UNVERSITYOF BEIRUT BEIRUT,LEBANOX -
(14) T h e 3,4'-isomer is aqsumed t o behave this way f r n m t h e I>e ha\-ior of t h e o t h e r isomer,.
