3958
E.A. BURNSA N D DAVIDN. H u m
unsaturated hydrocarbon is produced in a yield in excess of the saturate. Hydrogen and small amounts of other products are also formed. It is evident that, while a large fraction of the decomposition does proceed in a manner analogous to that of the photolysis, there are other reactions of varying importance that do contribute. One must be careful not to overgeneralize the similarities that do exist. The predominant effect of the radiolysis of the alkyl iodides is the rupture of the (2-1 bond. This occurs in spite of the fact that an appreciable fraction of the energy is initially absorbed by the hydrocarbon part of the molecule. It appears, therefore, that absorbed energy is rapidly distributed over the whole molecule and results i n primary chemical dissociatior~swhich are characteristic of the activated species. I n the present case preferential rupture is expected to occur upon localization oi‘ energj’ in the weaker carbon-iodine bond. Exclusive rupture of this bond in the alkyl iodides is of course known for low energy quanta from spectroscopic18 and photochemical8 considerations. In a similar manner organic acids are known to undergo preferential decarboxylation. l 9 (18) C,. Herzberg and C . Scheibe, Z. p h y s i k . C h e w , B I , 390 (1930). ( 1 9 ) I. A . Breger, J . Phys. Colloid C h e i i z , 52, 551 (1948).
[COSTRIBUTIOS FROM THE L)EPARTMEST
\.‘ol. 7s
lLIcDonell and Newton2” have found that in the high energy irradiation of the aliphatic alcohols the energy is directed to the a carbon atom. There is one other very important aspect to the present work. McCauley, Williams and Hamill have reported that in the photolysis of n-propyl iodide at 3261 A. there is very little isomerization (- 3%) of the n-propyl Zadical. Under more energetic conditions a t 2337 A. there is observed an appreciable yield of isomer (- %yo). In the present work isomerization amounts to only about 456. This very strongly indicates that a large majority of the radicals result from processes in which barely sufficient energy is given to the molecule to rupture the C-I bond. This contrasts with the assumption frequently made that since ioi1iziIlg radiation is very energetic, a very drastic effect upon the system is produced. As is of course el+ denced by the formation of minor products, some reactions do occur which involve a considerable amount of energy. In a large majority of the cases, however, the processes represent a miiiiniuni of energy in the reacting system. (20) W. R. McDonell and A. S. Newton, THISJ W R N A I . , 1 6 , 4Ii.51 (1954).
BUFFALO8, h T ~ YORK w
O F CHEMISTRY ASD LABORATORY FOR ISSTITUTE O F TECHNOLOGY]
SUCLEAR SCIESCE, hlASSACHUSETTS
Acetato Complexes of Lead in Aqueous Solution BY EUGENEA.BURNS’AND DAVIDN. HUME RECEIVED FEBRUARY 17, 1956 T h e existence of three complex ions, Pb(OAci+, Pb(OAc)*and Pb(OAcjs- has been shown in perchlorate solutions of lead containing excess sodium acetate and acetic acid. By the use of polarographic, potentiometric and solubility techniques successive “over-all” formation constants have been evaluated at a n ionic strength of 1.98 M. T h e average values of the “over-all” formation constants obtained by the three techniques for the mono-, di- and tri-acetato complexes are 145, 810 and 2950, respectively. The presence of anionic species was confirmed b y means of an anion-exchange technique.
The formation of complex ions by lead salts in acetate media has been investigated in the past by several authors. 2--x The striking disagreement among them both as to the number and stability of the complexes formed is illustrated by the results summarized in Table I. In almost every instance there has been some evident source of potential error: the use of nitrate media, which have recently been shown to cause complex formation with variable ionic strength or standard state, ( I ) Abstracted from a thesis submitted bi- Eugene A . Rrirns i n par^ tial fulfillment of t h e requirements for the degree of Doctor of PhilosoIihy from t h e Massachusetts Institute of Technology, 1950. ( 2 ) A . Jacques, 7’rons. Fni’nday .S’oc., 6 , 22.j (1909). 1 3 ) S. 11. ltdmonds and S . Birnbarim, T H I SJoufissi., 6.2, 23Gi t1NU). 1 +j B C . I’urkal-ashta and K . N. Sen-Sarma, .I. I ?id.l h e i i i . S‘UL., 23, 3 1 (194fi). ( 5 ) N. R . Das, S. Aditya and B. Prasad, i b i d , 29, Ifi9 (1952). ( f j ) S. Aditya and B. Prasad, i b i d . , SO, 213 (1953). ( 7 ) V. F. Toropova and F. hl. Batyrshina, Z h u r . A n a l . K h i m . , 4, 3 3 i
(1949). ( 8 ) S . Suziiki, .I. C h m . Soc. Japaiz, Piwe Chew?. .Srrlinn, 1 4 , 531 f I!I.i3) .
I I 11 I ~ e n I ~ c n ~A Ic. ~I Cn hriiitli i l i i ~ l 1 ) S.I I i i i r i r , ‘ f i l i i J O I ~ K 1 5 , 507 ( 1 < 1 5 3 J . l < c ~ l ~ i n > 'l'he\i\, Sl.t.> I n \ t , i f ' f c c h n i , l < i L y , l!l.jIi,
3960
E. A. BURNSAND DAVIDN. HUME
Vol. 78
standardized by potentiometric titration with perchloric acid. The ammonium salt solutions were standardized by the Kjeldahl method. Stock sodium perchlorate solutions were standardized gravimetrically by precipitation with potassium chloride. Iodate was determined iodometrically by means of the starch end-point using thiosulfate t h a t had been standardized vs. iodate.
mates the analytical concentration when an excess is present). Initial experiments when no acetic acid was present gave irreversible waves in which case the above method is not applicable. It was felt that in these experiments a competition for the lead, between hydroxide and acetate ions (the pH range was Results 3.22-8.06), was responsible for the irreversibility. Polarographic.-The polarographic results were To obviate this effect the solutions were made 1.96 interpreted by the method of DeFord and Hume.15 A ! in acetic acid, which in turn caused a junction This method consists of the graphical differentiation potential effect to appear in the comparison of soluof the Fo(X) term which can be obtained from po- tions with and without acetic acid, I t was also larographic data noticed that the diffusion coefficient of the lead ion was not the same in solutions with and without aceR ( X ) = antilog [ ( E I / : )-~ ( E I / ~ ) J log tic acid present. Table I1 shows the results of a (1) study made on the effect of acetic acid on this sysat 25.0°, where n is the number of faradays of tem. The shift of the half-wave potential toward electricity required per mole of electrode reaction, more negative values a t higher PH when acetate is El/, refers to the half-wave potential, I is the dif- present appears to indicate the presence of hydroxfusion current constant and the subscripts s and c ide competition. An even more striking effect was refer to the simple (aquo) ion and the same ion in evident in the absence of acetate as would be expected complexing media, respectively. I n solutions with from the known tendency of lead ions to hydrolyze excess ligand present and a constant ionic strength, in the absence of complex forming anions.16 The diffusion current constant and half-wave i t is known that potential in 1.96 144 acetic acid were plotted as a K ( X ) = 1 + KiX K2X2 R3X3 . . . Kn X” (2) function of acetate concentration and extrapolated where K1 . . . K, are the different “over-all” con- to zero acetate concentration and these values were centration constants and X is the equilibrium con- used for I, and ( E I / ~in) ~ eq. 1, respectively. I n centration of the ligand (which closely approxi- Table 111, experimentally determined half-wave potentials and diffusion current constants are given, TABLE I1 together with the calculated values of F(X).When EFFECT OF ACETICACID O N THE DIFFUSION CURRENT AND these were plotted against acetate ion concentraHALF-WAVE POTENTIAL tion, F2(X) and F3(X) were found to be straight NaOAc, HOAc, El/,# lines, demonstrating the existence of PbOAc+, PbpH Io v. vs. SCE M M ( O A C )and ~ Pb(OAlc),- in this range of acetate con0.493 0.0100 6.42 3.88 -0.4475 centrations with clear evidence against the existence ,0199 6.61 3.75 ,4476 of any significant amount of higher complexes. ,1078 5.39 3.71 ,4474 The extrapolation of the F(X)curves to zero ace,499 3.71 3.63 ,4460 tate ion concentration gave values for the successive ,988 4.43 3.48 ,4455 “over-all” formation constants which were 130, 840 1.966 3.25 ,4431 4.05 and 3000 for K1,K 2and KB,respectively. 0.0000 .000* 1.93 4.62 ,3686 Potentiometric.-Employing Leden’s” method, 0.0028‘ 1.954 1.67 3.92 .3644 the potentiometric determination of successive a Diffusion current constant in ma. l./mole mg.’/r sec.-’/z. * 0.0100 M HC104 present in solution. c Calculated from formation constants was performed a t an ionic dissociation of acetic acid. strength of 1.98 M . In this method
)
+
+
+
+
21
+
TABLE I11 ANALYSISOF Ell%OF LEADIN ACETATE MEDIUMWITH p = 1.98
where Fo(X) and n have the same meaning as before, and AE is the potential difference of concentration cells in which one half-cell contains no ace0,0000 -0.363ob 3.99* (1.00) . .. .002Sc .3644 3.92 .. . .. . , . . . tate and the other a varying known concentration .0394 .3895 3.66 8.59 193 of acetate. As in the polarographic solutions, 1.97 ,0893 ,4029 3.55 25.1 270 iM acetic acid was present in all the solutions, and .1896 ,4176 3.40 82.4 429 1.47 3.32 1.72 3.03 thus there was not a solution that contained abso,290 ,4282 3.37 189.4 649 968 2.09 3.19 .391 ,4368 3.29 379 lutely no acetate. To take this into account, the ,493 ,4431 3.25 627 1268 2.28 2.88 potential differences were extrapolated to zero ace2065 2.75 2.75 ,695 .4534 3.16 1436 tate ion concentration, the extrapolated AE being .997 .4655 3.03 3850 3880 3.72 2.89 1.424 ,4772 2.79 10360 7270 5.00 2.92 -1.8 mv.; then A E in eq. 3 was equal to A E ob1.901 ,4875 2.66 24350 12660 6.66 3.06 served 1.8 (in mv.). KI = K, = KI Table IV gives the experimentally determined 150 840 3000 potential differences corrected to zero acetate cona Diffusion current constant in ma. l./molemg.2/~sec.-1/p. centration together with the calculated values of Extrapolated value. Calculated from dissociation of acetic acid. F ( X ) . Once again, only three complexes were Na0.4~
,M
V.
ra
F.(x)
F~(X)
FdX)
FdX)
x io-’ x 1 0 - 8 .. .. .. .. 1.09 _. 1.34 ..
-
(15) D. D. DeFord and D. N. Hume, T H I SJ O U R K A L , 73, 5321 (1951).
+
(16) G. Mattock, A d o Chem. Scond., 8, 777 (1954). (17) I. Leden, 2. physik. Chem., A188, 160 (1941).
ACETATOCOMPLEXES OF LEADIN AQUEOUSSOLUTION
Aug. 20, 1956
3961
Table V gives the experimentally determined solubilities as a function of the acetate concentration together with the calculated values of F ( X ) . TABLE IV Three complexes are again shown to be present, the POTENTIOMETRIC ANALYSISOF LEAD IN ACETATEMEDIA successive "over-all" formation constants K1, Kz WITH p = 1.98 and K3, being 130, 770, 2350, respectively. AlNaOAc, E," FdX) FdX) though they are slightly lower than the constants A4 mv. F~(x) F , ( x ) x 10-8 x 10-3 obtained by potentiometry and polarography, the 0 . 002Sh 1.8 .. .... ... .. .. agreement is quite good. .0394 27.9 8.78 197 1.07 .. shown to be present and the values obtained for K1,K z and Ks are 155,820 and 3350, respectively.
,0893 .1896 ,290 .391 .493 .695 .997 1.424 1.901
41.8 58.0 67.4 77.0 B.9 95.2 107.3 120.1 130.7
.. .
25.9 91.4 190 401 687 1656 4250 11500 26250
279 1.39 476 1.69 . 652 1.71 3.07 1024 2.22 3.58 1390 2.50 3.40 2380 3.20 3.42 4260 4.12 3.31 8080 5.57 3.33 13800 7.17 3.34 K1=155 Kz=820 Ka=3350 Cal5 Potentials referred to extrapolated zero acetate. culated from dissociation of acetic acid.
Solubility.-The fact that the general principles of Leden's method may also be applied to solubility data has been known and used by previous investig a t o r ~ . ~I*t ~is~not ~ generally known, however, that solubility data may be used to give F ( X ) functions of exactly the same form as in the potentiometric and polarographic methods. These functions may be derived readily for the general case of the solubility of the compound M P A ,in complexing media (ligand X). The following equilibria need to be considered
+
&Ip A, = JhI+" qA-(Pa/Q) I\.I+a + X-r = MXCa-v
+ 2x7
M+a +
= %X-r =
(4) (5) (6) (7)
Mr;z+a-2r
M>;,+a-nr
TABLE V SOLUBILITY ASALYSISOF LEADIN ACETATEMEDIAWITH /I = 1.99 NaOAc,
M X 108, M
F~(x)
0.0000
0.604 1.912 3.01 5.40 7.92 14.05 22.2 35.5 60.8
(1.00) 10.0 24.8 79.9 172 54 1 1354 3450 10120
M
.0504 .lo07 ,2014 .3021 .504 .705 1.007 1.511
+
+
(18) E. L. King, THISJOURNAL, 71, 319 (1949). (19) G. C. B. Cave and D. N. Hume, ibid., T S , 2893 (1955). (20) E. L.Zebroski, H. W. Alter and F. K. Heumann, ibid., 73,5040 (1951).
Fx(X)
x
10-8
Fa(X)
x io-:
..
...
..
179 236 392 566 1070 1929 3423 6700
0.913 1.02 1.29 1.43 1.86 2.53 3.27 4.34
KI =
Kz
2.84 2.51 2 56 2 20 2.16 2.50 2.48 2.36 Ka = 2450
130
=
770
Discussion It has been shown by three independent methods that there are three acetato complexes, Pb(0Ac)+, Pb(0Ac)n and Pb(OAc)3- whose "over-all" formation constants are 145, 810 and 2930, respectively. Averaging the values of the three methods with equal weight was felt to be justified, for although the potentiometric and polarogTaphic determinations essentially measure the same quantity, the solubility determination is not necessarily as accurate. Figure 1 shows the distribution of lead in
In this case ( m / m ~ ) ( P f , ) /=P F o ( X ) where F o ( X ) has the same meaning as in the previous sections, and r n o and m are the solubilities of the compound in the absence and presence of complexing ligand, respectively. The slightly soluble compound chosen in this work was lead sulfate, and here there is also ,the possibility of hydrogen sulfate ion formation. It can be shown that in this case
where Ka is the ionization constant of the hydrogen sulfate ion a t an ionic strength of 1.99 and HOand H are the equilibrium hydrogen ion concentrations in the presence and absence of complexing ligand (acetate) , respectively. Since junction potentials would not be a factor in this case, the experiments with zero acetate also contained no acetic acid but instead were 1.00 millimolar in perchloric acid. Using a value of 0.084 for the dissociation constant of hydrogen sulfate ion a t an ionic strength of 1.99 M 2 0 and assuming the activity coefficient of hydrogen ion to be unity in this medium, it was found that Ka)/(H K a ) , is the correction factor, (Ho practically constant and is equal to 1.007 f 0.007, over the pH range of 3 to 5 which was used here.
F,(x)
/
Pb (OAc); PbOAc'
40
t
/ \
DhlAAr\r,
[log O A c - I
Fig. 1.-Percentage of lead in its various forms as a function of the free acetate ion concentration,
its various forms a t an ionic strength of 1.98 M as a function of the free acetate ion. That these values do not confirm those of earlier investigators is obvious, and the reasons are undoubtedly the same as were suggested in explanation of why previous workers disagreed so radically among themselves. The discovery that the carefullv Dlanned and executed work of Edmonds and Biinbaum, long ad-
3962
E. A. BURNSAND DAVIDN. HUME
Fig 2 -Thermometric titration curves with 0 990 M ammonium acetate Circles, titration of 50.0 ml. of 0 0600 J!lead nitrate; triangles, blank titration of 50 0 ml of water.
mired as a classic, had failed to reveal the higher complexes came as a surprise, however. The most likely explanation probably lies in the fact that lead ion is extensively hydrolyzed and polymerized as the PH is increased.16 Reaction of hydroxide ions with lead iodate may produce a solid phase in which the lead to iodate ratio is no longer 1: 2 . This would mean that the iodate concentration in solution could not be used to calculate the solubilities of lead iodate directly, as in their experiments. Some spot checks were made of the solubility of lead iodate in various solutions (Table VI), and it was evident that the lead to iodate ratio in solution differed appreciably from the theoretical in a number of instances. When it is remembered that the solubility values were cubed in the subsequent calculations, it is not unreasonable that the accumulated errors might obscure the presence of the relatively weak higher complexes. TABLE 1.1 ISVESTICATIOSS OF THE SOLUBILITY OF LEAD IODATE IN ACETATEMEDIA Soh equilibrated
fIo8-j
(Pb-')
\rater 1 .OO T . I SH4C104 I) 99 SH4O.ic
0 722 3 9,5 18 3
0 340
1 71 8 59
2 12 2 27 2 13
1 00 JI NaOAc, 1 0 98 Jf HOAc
13 4
6 67
2 02
x
1
10-4
x
10-4
([Or)/ (Pb+-)
The thermometric titration curve presented by Purkayastha and Sen-Sama4 showed a very sharp end-point a t the I : 1 complex which is not consistent with the results one calculates from the formation constant they gave ( K , = 2 5 ) , Le., at the endpoint, only 64Yc of the lead would be complexeti.
Vol. 78
We have repeated their work, titrating 50.00 ml. of 0.0600 M lead nitrate with 1.00 M ammonium acetate. Figure 2 shows the result of this titration. It is interesting that the end-point was quite rounded and the initial slope was endothermic, both observations being contrary to those of the Indian workers. One would not expect to see further changes of slope after the end-point since the higher complexes are not formed extensively at these concentrations, and the slope after the end-point should correspond to the exothermic heat of dilution of ammonium acetate as was observed in Fig. 2. The actual presence of an anionic complex in lead acetate mixtures was demonstrated by Fronaeus' ion-exchange method.21 Here the distribution coefficient of lead between resin and solution was measured polarographically as a function of the ligand concentration over the range of 0.02 to 1.4s Macetate ion. The PH was held constant by maintaining a 1: 1ratio of acetic acid and sodium acetate, and the ionic strength varied accordingly. In this method, the existence of negatively charged complexes is shown by the appearance of a maximum in the distribution coefficient us. ligand concentration function. Although the data are not highly accurate because of the limitations of the analytical method a t very low lead concentrations, the data of Table VI1 show clearly that such a maximum is present. The location of the maximum a t approxiTABLE VI1 ION-EXCHANGE STUDIESo s THE LEAD ACETATE Srsrear T-. earl ~n ...
_I
Initial lead
x
loa,
.V
NaOAc, M
Anion exchanger 2 00 0 020 2 00 040 2 00 080 2 00 480 2 00 980 2 00 1 180 Cation exchanger 2 40 0 050 2 40 1 500
Final lead x 106, M
378 107 20.5 3.0 3.7 9.4 380 2300
resin, mole/g.
x
106
Distribution CMK/CAl.
1 /fi.
8.2 9.5 9.9 10.0 10.0 9.95
0.022 0.088 0.48 3.5 2.7 1.1
10.1 0.5
0 ,026 0.00022
mately 0.5 M acetate ion corresponds closely to the acetate ion concentration where the average ligand number is 2 (Fig. a ) , z.e., the maximum concentration of the uncharged lead acetate molecule, as is predicted by Fronaeus' equations. The experi ments with a cation-exchange resin, although highly suggestive of the presence of anionic coni plexes, are not as conclusive as the anion-exchange results because the low distribution coefficient a t high acetate concentrations could be due t o the formation of the uncharged species. Acknowledgment.--This work has been supported in part by the U. s.Atomic Energy Commission. CAMBRIDGE, MASS (21) S. Fronaeu.;, C z ~ ? i s k K. r m . Ti,]., 65, 1 f l H R 3 j
