798
L. B. RYLAND, G. S. RONAYAND F. M. FOWKES
Vol. 62
for OH or by an increase of the chain length with a radical and for unequivalent coupling of protons in CH2 group. Again the specific radicals cannot be the same radical. assigned. In France Combrisson and Uebersfeld’ have inOrnithine, Lysine, Arginine and Citrulline.-In vestigated with electron spin resonance the effects the electron spin resonance curves for X-irradiated of atomic pile irradiation on a number of amino ornithine hydrochloride, NHz(CH2)&HNH2COOH- acids. They reported a triplet resonance for several (HCl), and lysine hydrochloride, NHz(CH2)4- amino acids and no observable signal for others. CHNHzCOOH(HCl), one again sees the qualita- We do not yet know the causes for the differences betive differences caused by an increase of one CH2 tween our observations and theirs, but they must group in the chain. See Fig. 16. (Compare with arise from differences in the experimental conditions asparagine and glutamine above.) Figure 16 also or in the methods of irradiation. gives the resonan.ces of X-irradiated arginine and (7) J. Combrisson and J. Uebersfeld, Cornpt. rend. Acad. Sci. citrulline. All these resonances are complex and (Paris), 268, 1397 (1954); J. Uebersfeld, thesis for degree of doctor give evidence for the existence of more than one of physical science, University of Paris, 1956.
EQUILIBRIA IN AQUEOUS SOLUTIONS OF COPPER(I1) CHELATES WITH U,u’-DIPYRIDYL, O-PHENANTHROLINEAND ETHYLENEDIAMINE1 BY L. B. RYLAND, G. S. RONAYAND F. M. FOWKES Shell Development Company, Emeryville, California Received December 29, 1067
The equilibria of species in solutions of the 1:1 chelates of cupric nitrate with a,a’-dipyrigl, 9-phenanthroline or ethylenediamine have been investigated by potentiometric titration with sodium. hydroxide. quilibria have been calculated from the sodium hydroxide consumption in solutions of different cupric chelate concentration a t constant p H . It was found that in solutions of the 1:1 chelate of cupric ion and dipyridyl (containing 0.1 M KNOl to maintain constant ionic strength), the system studied in greatest detail, three species are present in measurable quantities: an acidic species A, DipyCu(HzO)zf+; H an uncharged but water-soluble basic species Bz, .DipyCu(OH)zo; and a dimeric species (BJz, DipyCu(
0
)CuDipy
f
+.
0 H
Presumably a species B1, DipyCu(HzO)OH+, also exists but not in measurable concentrations. I n acidic solutions species A predominates, in neutral solutions species (BJz predominates, and in alkaline solutions species Bz predominates; species (B& increases in concentration at the expense of other s ecies with increase in the over-all concentration of copper chelate. (B& 2 H + 2H20, = 10-10.14, and for theequilibriumA *Bz 2H+, K A = 10-16.28. At 25” for the equilibrium 2A Analogous constants for solutions of the 1:l chelate of cupric ions and o-phenanthroline are Kaz = 10-1°-86 and K A = 10-17.0. I n both systems the constants increase with temperature, AHa2 is 8.2 and 8.3 kcal./mole and AHA is 5.4 and 4.2 kcal./mole for the dipyridyl and o-phenanthroline chelate, respecti)vely.
+
+
Zaz
Introduction The hydrolysis of hydrated metal ions has been studied in some detail by others2-6 who have shown that hydrated metal ions such as cupric, uranyl, bismuth and scandium hydrolyze to form basic ions and dimers or polymers of the basic ions. A similar investigation has been undertaken with hydrated chelated metal ions. The present investigation has been confined to the 1: 1 chelates of cupric ion with a,a’-dipyridyl or with 1,lO-o-phenanthroline which are of interest as catalysts in hydrolysis reacti0ns.l The object of the investigation is to identify the various species in such solutions and to determine the appropriate equilibrium constants as a function of concentra(1) (a) Presented a t the 128th Meeting of the American Chemical Society in Minneapolis, Minnesota, September, 1955; (b) this paper reports work done under contract with the Chemical Corps, U. S. Army, Washington 25, D. C. (2) K. J. Pedersen, #. Danske Videnskabernes Selskab, 2 0 , No. 7 (1943). (3) F. Graner and L. G. Sillen, Acta Chem. Scand., 1 , 631 (1947). (4) 8. Ahrland, ibid., 8 , 374 (1949). (5) M. Kilpatrick and L. Pokras, J . Electrochem. Soc., 100, 85 (1953). (6) S. Hietanen and L. G. Sillen, Acta Chem. Scand., 8 , 1607 (1954). (7) T. Wagner-Jauregg, el al., J . A m . Chcm. Soc., 77, 922 (1955).
+
tion, pH and temperature. A few results with ethylenediamine are appended. Experimental Details Materials.-The a&-dipyridyl was certified reagent grade from Fisher Scientific Company and its melting point indicated a purity of 99.9%. The 1 ,IO-o-phenanthroline monohydrate was manufactured by the G. Frederick Smith Chemical Company. The ethylenediamine monohydrate was obtained from Eastman Kodak Company (No. 1915) and had a purity of 97.2% by acid titration. The cupric nitrate and potassium nitrate were standard reagent grade chemicals. Solutions for titration were maintained free of COZ and contained 0.1 M KNOa and 4.75% of propylene glycol. Apparatus.-Potentiometric titrations were made with a Precision-Shell Dual AC Titrometer, using a jacketed titration beaker through which thermostated water was passed to give temperature control.
Discussion of Results Equilibria of Species in Solutions of 1 : l Cupric Dipyridyl Nitrate.-The stability of the 1 : 1 chelate is sufficient* that it is treated as an undissociable entity. The titration of this chelate with sodium hydroxide (Fig. 1) shows that it loses two hydrogen ions in two steps. Designating the minimum (8) Cf.A. E. Martell and M. Calvin, ”Chemistry of the Metal Chelate Compounds,” PrenticeHall Publ., Inc., New York, N.Y., 1952.
July, 1958
799
EQUILIBRIA OF Cu(I1) CHELATES WITH Q,Q'-DIPYRIDYL
hydration necessary to explain the observed reactions, we write
+
DipyCu(H20)2++
DipyCu(H20)OH+ H + DipyCu(OH)20
Conditions: Titrating medium, 0.1 M KNO, solution containing 4 . 8 % ~
+ 2H+
which we abbreviate to A ~ B ~ + H + ~ B z + 2 H +
9 -
giving rise to equilibrium constants The family of curves in Fig. 1 cross at one mole of sodium hydroxide per mole of chelate, which occurs at pH 8.1. At this ratio [A] = [Bz],so from the definition of K Awe see that the pH of the "crossover" is p K ~ / 2 . The pH of the crossover is independent of concentration, but a t any pH below or above it the titration curves shift with increase in concentration toward an hydroxyl to copper ratio of 1.0. This means that species B1 forms a dimer or polymer. The formation of polymers of B1 can be written as nA
g
8-
7-
+ n H + + nHzO
(BJn
The equilibrium constants have been determined from the titration data (in the pH range of 6-7.5 where [Bz] is small compared with [A]) in the followingmanner. For any point in a ,.itration curve let Y eaual the total eauivalents Der liter of hvdroxide ionsumed up to chat point
I
6
Pi1
+ 2Pzl +
I
OH-/CU++.
n[(BdnI
10-2
1
(1)
I
n=2
From this series we may determine KA, and the K,, values if we know [A] as a function of pH. The value [A] is calculated from the initial concentration of copper chelate [A]i
-
Y
+ [Bzl
(3)
1' - 2[Bz1 =
Y
+ [&I)
P
-
i ,,
-
u F
10-4
10-5 10'
102
[AI/[Htl
+
/
\
u
I
IN-
and ~ K isA obtained by doubling the pH of the "crossover." Substitution of (3) into (2) gives
-
d
0 pH6.5 D DH 6 . 7 5
-2
in which [Bz] is closely approximated by
KA, ( [Ali
I
moles/mole.
m
Substitution of equilibrium constants gives
[A] = [Ali
I
Fig. 1.-Titration of 1:1 dipyridyl cupric nitrate a t various concentrations with sodium hydroxide.
A
Y
//
10' = (Ai-Y)/[Ht].
Fig. 2.-Hydrolysis of the 1:l cupric dipyridyl nitrate a t 25' pH 6.5 to 7.0, where "A" represents DipyCu ( H20)z$+, and B1 DipyCu( H 2 0 ) 0 H+.
By plotting the logarithm of Y - 2[Bz] vs. the Iogarithm of ([Ali - Y [BJ) as in Fig. 2 the resulting line has a slope of E , the average number of BI units in the aggregating species, and from the coordinates of a point in a section of given n, Ka, may be calculated from
+
In several cases [Bz] turned out to beynegligible and log Y os. log ([Ali - Y)is plotted. At pH 9 or above, where [Bz] is large and [A] very small, we may use a similar treatment by considering addition of acid to a solution of species Bz nBz
(5)
+ nH+
Kgn =
[(Blhl [Bzln[H+ln
L. B. RYLAND, G. S. RONAYAND F..M. F O W K E ~
800
Vol. 62
termined at pH 6.5, 6.75 and 7.0 for all eleven concentrations (these are listed in Table Is). The log Y us. log ( [ A l i - Y)/[H+] relation derived from equation 4 is shown in Fig. 2. A constant slope of 2.0 over the whole range of concentration shows that (Bl)z is the only form of B1 present in measurable concentrations. This may be represented as
0 pH 9 . 0 0 pH 9 . 2 5 ApH9.5
H 0 DipyCu( )CuDipy
++
0 H
10-5
10-2'
IO-"
10-l)
[B1][Ht]
(rnolen/liter)'.
= (U-[A]l)[Ht],
Fig. 3.-Hydrolysis of the 1 :1 cupric dipyridyl nitrate at 25", pH 9.0 to 9.5, where BZrepresents DipyCu(0H)z.
By using the above expression in equation 1 we obtain Y (total hydroxide consumed) as a function of [Bz] Y
=:
+5
~ [ B z.t ] [Bzl[H+l ~
KA*
nKfin[Bz]"[H+]" (6)
n=2
[Bz]is determined from experimental measurements by [BzI = Y
- [Ali + [AI
in which [A] is usually negligibly small, but can be closely approximated by [A]
(2 [ A i - Y)l'z
[H+]
a2
In terms of experimentally determined quantities, equation 6 becomes @[A],
- 2[A] - Y ) = ( Y + [A] - [Ali) [H+l +
Ka2 = 6.5 X 10-12 =
m
C fiK~n(Y+ [AI - [AIi)"[H+l"
(7)
n=2
This relation is used for determination of fi and KBnby plotting log (2[A]i - 2[A] - Y )us. log (Y [A] - [Ali) [ H + ] , as in Fig, 3. The slope gives f i and the coordinates give KBnat any p H or concentration, for
+
log Kgn = log (2[AIi
- 2[Al - Y ) n n log (Y
+ [AI - [Ali)[H+l
(8)
Determination of Ranand Kpnallows calculation of KA (provided ii is the same in both p H regions), for _ = - = KnKA'n
KBn
KJKA:"
(KA,KA*)I= K A ~
Furthermore, in the region of higher p H , the plot of titration data according to equation 7 in Fig. 3 also shows a constant slope of 2.0 and shows that in this p H region, also, the only form of B1 is the dimer (B&. The presence of other species should be evident in Figs. 2 and 3 by curvature at either end of the straight lines to higher values of Zn [(Bl),J. The presence of species B1 would curve the lower end of the line upward and the presence of (B1)l or higher polymers would curve the upper end upward. Of course, some B1 must be present, but the maximum value of K A consistent ~ with the data of Fig. 2 is about lo+.'. The equilibrium constant Kaz is found by equation 5 to be 1.8 X 10-l1 or 10-10.74,K p z is found by equation 8 to be 10+21.83, and K A is found by equation (9) to be 5.2 X 1O-l' or 10-16.28 The p H of the crossover should be p K ~ / 2or 8.14; the observed crossover is actually $H 8.1. Eguilibria at 0 and 41.2°.-Titrations were made a t 0 and 41.2' with 5 X 10-3 and 5 X mole/ liter of copper chelate (Fig. 49. Values of Y were determined a t p H 6.25, 6.50, 6.75, 7.0, 7.25, 8.5, 8.75, 9.0 and 9.25 (Table IIg). In calculating Y , the free hydroxyl ion concentration at 0 and 41.2' was calculated from the p H measurement and using 10-14.95and 10-13J, respectively, for the dissociation constant of water. The graph of log Y us. log ([AJ - Y ) / [ H + a] t O'gives a straight line of slope 1.8 to 2.0 (6)and it is found that
(9)
This is especially useful in these systems where [B1] is immeasurably small so that Kn, KA, and KA are indeterminable, but K,nl Kfin and K A can be determined accurately. Equilibria at 2S0.-From the titration curves shown in Fig. 1, values of Y and [A]i - Y were de-
10-11.19
For the data at p H 8.75 - 9.25, log (2[A]i - 2[A] Y )was plotted vs. log ( Y - [A]i [A]). A straight were line of slope 1.9 to 2.0 and a KBZof 10+21.86 obtained. By equation 9
+
so the ~ K = A 16.52. This checks with the ~ K A of 16.5 obtained by doubling the p H of the crossover ( p H 8.25). At 41.2", we obtained Z = 1.85, and K,z = 10-10.32. This 6 is considered to be within experimental error of 2.0. The data obtained a t p H 8.5, 8.75, 9.0 and 9.25 when plotted according to equation 7 give KBz = 10+21.62and KA = 1O-l6 92. The PKA of 15.92 checks the value of 15.90 obtained by doubling 7.95, the p H of the crossover. Temperature Coefficients of Kaz and KA.-When the values of log Kaz and log K A are plotted us. (9) A t the request of the Editor, Tables 1-111 and Figs. 4-6 are omitted but are available from the authors, or as photoprints or microfilm (Dooument No. 5520) for $1.25 from the AD1 Auxiliary Publicatiom Project, Photoduplication Service, Library of Congress, Washington 25, D. C.
July, 1958
HYDROGEN OVERPOTENTIAL ON ELECTROPLATED COPPER-TIN ALLOYS
1/T the three points do not give a straight line, but rather a curve with increasing slope a t higher temperatures. The slopes at 25" give AHa2 = 8.2 kcal/mole, and AHA= 5.4 kcal/mole. Equilibria of Species in Solutions of Copper oPhenanthroline (1:1) Nitrate.-As in the case of copper dipyridyl, several species of copper ophenaiithroline exist in aqueous solutions A-o-PhCu(HiO)z + + Bi-o-PhCu(Hz0) (OH) + Bz-0-PhCu (0H)p (BI)Z-(O-P~CUOH) 2 ++
influenced by the basicity of the ligand; as the basicity (or p K a ) of ligands increases, the acidity of the chelates should decrease. Thus in comparing dipyridyl (pKa= 4.4) with o-phenanthroline (pK, 5.0) as ligands for cupric ion, it is found that the o-phenanthroline chelate is the weaker:acid. TABLEIV EQUILIBRIUM CONSTANTS FOR COPPERDIPYRIDYLA N D COPPER0-PHENANTHROLINE IN 0.1 M KNO, Constant
Kaz
The equilibria between these species are as shown previously. The titration curves at 25 and 41.2' with 5 and 10 X lo-* moles/liter of chelate are shown in Fig. 5,9 the values of Y at pH 6.5, 6.75, 7.0, 7.25 and 7.5 are shown in Table IIIj9and the graphs of log Y vs. log ([Ali - Y ) are shown in Fig. 6.9 At 25", % is found to be 2.0 throughout the whole range of pH values and concentrations, and Ka2 = 1.27 X lo-" = 10-10.90. The pH of the crossover is 8.50 so K A = 10-17.00. At 41.2", % is also 2.0 for the whole range of pH values and concentrations, and Kea = 5 X lo-" = 10-10.30. The pH of the cross-over is 8.35; thus, KA = From the values of K,2 and KA at 25 and 41.2", one may calculate AHoz = 8.3 kcal./mole, and AHA = 4.2 kcal./mole. Comparison of Equilibrium Constants.-In Table IV are shown the equilibrium constants at various temperatures for copper dipyridyl (1: 1) nitrate and copper o-phenanthroline (1: 1) nitrate. The tendency to dimerize to species (B& is about the same, but the o-phenanthroline chelate is definitely the weaker acid as evidenced by the lower value of KA. The acidity of 1: 1 metal chelates undoubtedly is
801
KA
Tyw.,
C.
0 25 41.2 0 25 41.2
Copper dipyridyl
0.65 x 1.8 X 4.8 x 3.0 x 5.2 x 12.0 x
10-1' lo-" 10-11 10-17 10-17 10-17
Copper o-phenanthroline
. . . ,......
1.27 x IO-" 5 x 10-11
.... . . . . . . 1.0
2.0
x x
10-17 10-17
Equilibria of Species in Solutions of 1 :1 Cupric Ethylenediamine Nitrate.-In all of the foregoing work the solutions contained 0.1 M sodium nittrate to keep activity coefficients of the metal chelate ions constant over wide ranges of concentration. However, in the case of the ethylenediamine chelate, addition of salt caused formation of a coppercontaining precipitate so in this system salt could not be used and activity coefficients were not constant. Plots of the data as in Figs. 2, 3 and 6 have a very wide scatter of points. Nevertheless, the titration curves appear similar, with a cross-over at pH 8.6 (indicating a ~ K ofA17.2) and the value of Kaz is estimated to be about 10-ll,e. Acknowledgment.-The authors wish t o thank Dr. J. N. Wilson for advice and encouragement and Mrs. Helen B. Heppe and Mr. Sherwood C. Beckley for assistance in obtaining precise titration curves.
HYDROGEN OVERPOTENTIAL ON ELECTROPLATED COPPER-TIN ALLOYS BY I. A. AMMARAND H. SABRY Department of Chenzistry, Faculty of Science, University of Cairo, Cairo, E g y p t Received Auoust SO, 1957
Hydrogen overpotential has been measured on electroplated Cu, Sn and Cu-Sn alloys, in 1.0 N HC1 at 30". Six alloys ranging in composition from 12 to SO$'& Sn have been studied. The overpotential has been measured in the current density range 10+ to 3 x 10-8 amp./cm2. The results have been analyzed statistically and the 95% confidence limits, for the mean overpotential values, have been calculated. Attempts have been made to esplain the overpotential-alloy composition relation through the dependence of overpotential on the heat of adsorption of hydrogen on the alloy.
Introduction Overpoteiitial studies on alloys are numerous. Thus Newbery' observed that the overpotential 7 was higher for Pb-Hg and Zn-Hg alloys than the corresponding values for the components. Fischer2 measured 7 on a number of alloys and found it to be independent of the composition of the alloy and equal to that of the component with the lower overpotential. Harkins and Adams3 found the over(1) E. Newbery, J. Chem. SOC.,109, 1051 (1916). ( 2 ) P. Fischer, Z . p h y d . Chem., 113, 326 (1924).
(3) W. Harkins and H. Adams, THISJOURNAL, 29, 205
(1926).
potential on monel metal t o be lower than the corresponding values for the components. For Cu-Ni alloys, Raeder and co-workers4Jobserved a gradual change from the overpotential on Cu to that on Ni. Sharp maxima and a flat minimum mere, however, observed for the other alloys studied by the above auth0rs.~#5DeKay Thompson6 studied the relation between the overpotential and composition of brass, and observed a sharp minimum between 15 (4) M. Raeder and J. Brun, 2. physik. Chem., 153, 15 (1928). ( 5 ) M. Raeder and D. Efjestad, ibid., 8140, 124 (1928). (G) M. DeKay Thompson, Trans. Electrochem. Soc., 69, 115
(1931).
