5782
JAMES
VOl. 75
I. WATTERS,JOHN G. MASONAND 0. E. SCHUPP, TI1
the equilibrium constant, per mole of diborane, is loglo Kip (atm) = 7.478
- 6205/T
The measured kinetic constants which were cited by the various authors in the form of ( k z K ; t )
now can be evaluated separately. These have been assembled in Table I for ready reference, expressed in the form loA exp(-E/RT), with concentration units in [moles cc.-l]. ITHACA,
N. Y .
[CONTRIRUTION FROM MCPHERSON CHEMICAL LABORATORY, THEOHIO
STATE
UNIVERSITY]
Investigation of the Complexes of Mercury(I1) with Ethylenediarninetetraacetic Acid Using the Mercury Electrode BY JAMES I. WATTERS,JOHN G. MASON'AND 0. E. SCHUPP, I11 RECEIVED JULY 2, 1956 Mercury( 11) forms complexes with ethylenediaminetetraacetate ion which are reversibly reduced a t the dropping mercury electrode in the acidic pH range and to a p H of 10 if a relatively large excess of the ligand is present. The composite anodic-cathodic polarographic waves with excess ligand have slopes of 0.030 volt corresponding to a reversible 2 electron transfer. The relative magnitudes of the anodic and cathodic diffusion currents as well as the 0.030 volt shift in the wave for a tenfold change in the ethylenediaminetetraacetic acid concentration prove the complexes contain one mercury( 11) and one ethylenediaminetetraacetate ion. The mercury pool electrode rapidly reaches equilibrium with the complexes in solutions containing excess ligand from a p H of 10 t o at least 2. The potential of the electrode as a function of pH and ligand concentration has permitted the interpretation of the equilibrium and the evaluation of the equilibrium constants. Below a pH of 3.5, the association of one hydrogen ion by the bound ligand occurs while above a pH of 8, the association of one hydroxyl ion occurs. The predominant species in the PH range of 4 to 8 is Hg(enta)2-, with KIOO = 10z1.64 * a.o2 a t 25' and p = 0.1. I n more acidic solutions Hg(Henta)- also occurs with K110 = 1014.62 * 0.I'J. In more basic solutions the second species is Hg(OH)(enta)8- with KIOI= 1 O Z 6 . ~ 0.16.
Introduction After the completion of the polarographic portion of the present investigation a similar study was reported by Matyska and Kossler.2 I n agreement with our results they found three complexes Hg(enta)Z-, Hg(Henta)- and Hg(OH)(enta)3where enta indicates the ethylenediaminetetraacetate ion. However, their complexity constants were consistently smaller than ours. It was concluded that the difference was caused by an unreliable saturated mercury(1) sulfate reference electrode to which they assigned a potential of 0.417 volt instead of the theoretical 0.65 volt. More recently Goffart, Michel and Duychaerts3 as well as Schwarzenbach, Gut and Anderegg4 have studied the complex polarographically. The former authors studied the system by the conventional polarographic method in the pH range of 4.5 to 9.2. Their data indicated a single complex Hg(enta)2in this pH range with a complexity constant of 1032.16 a t 25' and p = 0.065 0.02. Schwarzenbach, Gut and Anderegg4 obtained polarograms of a solution containing two metal ions and insufficient enta to completely complex them. From the wave height of one metal ion in the uncomplexed form the known complexity constant of one of the complexes, and the known total concentrations, they calculated the concentrations of all metal species and finally the unknown equilibrium constant. I n this way the complexity constants for a large num(1) Presented in part before the Division of Physical and Inorganic Chemistry, American Chemical Society, a t Kansas City, Mo., April 1, 1954. Taken in part from a thesis by J. G. Mason submitted in partial fulfillment of the requirements for the Ph.D. degree, The Ohio State University, 1955. (2) B. Matyska and I. Kossler, Collection Czech. Ckem. Comman., 16, 221 (1951). (3) J. Goffart, G. Michel and G. Duychaerts, Anal. C h i n . Acta, 9, 184 (1953). (4) G. Schwarienbnch, K Gut and G. Anderegg, Helv. Chim. Acta, 37, 939 (1954).
ber of cations were calculated. Although the method did not succeed for mercury(II), they included equilibrium constants presumably calculated on the basis of the shift of half wave potential. For the simple and the acid complexes of mercury(11) a t p = 0.1 and 20" their values, Kloo = 1021.80 and Kilo = 1Ol4J4, are in good agreement with ours. Theoretical The general mathematical procedure used in the present study was presented in an earlier paper.6 Since all of the complexes contained mercury(I1) and the ligand in a (1:1) ratio, the calculations were relatively simple. The various equilibria involving the association of hydrogen ion with the bound ligand were first considered in terms of the concentration of ethylenediaminetetraacetate ion and the activities of hydrogen ions, instead of ions such as Henta3- and Hzenta2-, since this greatly simplified the calculations. These constants, indicated by primes, were finally converted to the desired forms. I n the following equations, parentheses indicate concentrations while brackets indicate activities. Hge+
+ enta4-
Hg2+ f s H +
Hg(enta)a-; ( Hg(enta)2-) Kioo = ( Hg2+)(enta4-)
(1)
+ enta4- I_Hg(H.enta)(Z-*)-; K'iso
(Hg( HI enta)(enta4-) (Hg2+)[H+la
H g 2 + f 10H- -I- enta4- I_ Hg(OH)t enta
(2)
(2+t)-;
Kiot = (Hg( 0H)tenta ( 1 + t ) - ) (Hg")[OH-] "enta4-)
(3)
I n terms of these equilibria, eq. 8 for Leden's6 (5) J. I. Watters and John G. Mason, THIS JOURNAL, 78, 285 (1956). (6) I. Leden, 2. pkysik. Ckem., 1888, 160 (1941).
MERCURY COMPLEXES WITH ETHYLENEDIAMINETETRAACETIC ACID
Nov. 20, 1956
function, Fo,of the previous papeld becomes
(v)) -
( [ H + ] + K ’ l ~ ~ ( e n t a 4[H+] + K’1lo(enta4-) -) [ H + ] *+ . . . 4-K101(enta4-) [OH-] + Klo2(enta4-) [OH-I2 + . . . (4)
~,(enta‘-, H + , OH-) =
antilog
C ,
1=
K1oo(enta4-)
The following expressions of Leden’s equations which are obtained readily by the subtraction of KIOO and division by [H+] or [OH-] proved to be the most useful.
The function Fl is useful since its value remains constant in the range where Hg(enta)2- predominates. The functions Fzaand F 2 b are useful since they remain constant in the acidic and basic ranges where Hg(Henta)- and HgOH(enta)3-, respectively, predominate. Functions F1, FZa and F 2 b were adequate for evaluating the complexity constants for the three complexes found. The constant for the acid complex was converted to that for the equilibrium
5783
Experimental
The potentiometric and polarographic apparatus, the general experimental method and the preparation of the mercury(I1) solution were described in a previous publication.’ The cell was separated from the saturated calomel electrode by a salt bridge containing 0.1 M KNOs in 3% agar agar. All experiments were performed at 25 2c 0.1’ using solutions containing 0.1 M KNOI to produce an ionic strength of 0.10 f 0.01. The ethylenediaminetetraacetic acid solutions were prepared on a weight basis from C.P. grade acid, Versenes Inc. The molarity was verified by titrating with standard HCl using the glass electrode. Solutions of intermediate PH were prepared by mixing in various ratios the acidic and basic solutions having the same concentration of mercury( 11) and ethylenediaminetetraacetic acid. Potassium hydroxide and freshly boiled nitric acid were used to adjust the PH of the acidic and basic stock solutions. The capillary used for determining the polarographic diffusion current constant had a drop time of 5.89 sec. and delivered 1.278 mg. of mercury per second a t the potential of the S.C.E. corresponding to a capillary characteristic, rno’/ato’/fi, of 1.664 mg.*/r sec.’/i.
Results and Discussion Reversible anodic-cathodic polarograms were obtained for solutions containing mercury(I1) ion with a small excess of enta using 0.05 M NazBa07 a t p H 8.5 as the supporting electrolyte with 0.005% gelatin added as a maximum suppressor. The anodic-cathodic polarogram of a solution 0.5 X M i n Hg2+and 1.0 X low3Af in enta furnished H g 2 + I l e n t a 3 - z Hg(Henta)-; considerable information concerning the nature of the complex. Since the anodic and cathodic diffusion currents were approximately equal it was evi(Hg(Henta)-)K4 dent that the complex must contain the mercuryKilo = ( Hg2+)(H +)(enta4-) = K’llO x K4 (9) (11) and the enta in a (1:l) ratio. The excellent where k - 4 is the fourth dissociation constant of continuity of the wave across the zero current axis indicated a reversible electrode reaction. I n the ethylenediaminetetraacetic acid. Schwarzenbach expresses the equilibrium for the absence of mercury in the solution well shaped anodic waves with similar half wave potentials acidic complex in eq. 8 thus were obtained. With the same concentration of Hg enta2H + If Hg HentaM but with the mercury(I1) ion enta, 1.0 X (Hg Henta-) absent the anodic diffusion current was, as ex= K’llo/Kloo (10) Ka = (Hg enta2-)[H+] pected, twice as large as in the former polarogram. The value of AE/log(id, - i) in both the anodicwhich is the reciprocal of the acidity constant of Hg cathodic and the anodic waves was 0.032 v., provHenta-. For the alkaline equilibrium he writes ing the electrode reaction to be essentially a reversHg2+ enta4H i 0 3 Hg(0H) enta8H+ ible two-electron transfer. Similar waves were obtained in more alkaline solutions and below a p H of 4. However, approaching a pH of 7 the slope For this conversion and for the calculation of tetra- increased to more than double this value. Similar acetate ion concentration, the acidity constants results were obtained in unbuffered solutions 1 M determined by Schwarzenbach and Ackermann? in KNO, in which the PH was adjusted with nitric were used. For the successive hydrogen ion dis- acid or potassium hydroxide. When the enta concentration was increased to sociations a t 20’ in 0.1 N KCl they reported pK1 0.1 M the wave became continuous across the zero = 1.996, pKz = 2.672, pK3 = 6.161 and pK4 = 10.262. Since one ligand molecule is complexed current axis. There was no constant anodic difthroughout the range investigated, the concen- fusion current region due to the large enta concentration. The slope of 0.030 v., obtained throughtration of completely dissociated enta ion is out the fiH range of 2 to 12, indicates a reversible (entad-) = 2-electron transfer. Another advantage of the excess ligand concentration is the buffering capacity in the absence of foreign anions other than nitrate. I n 0.1 MKNO3 the diffusion current constant of the where C1 and C, are the total concentrations of Hg enta2- complex was 3.09 microamperes milliethylenediaminetetraacetic acid and mercury(I1) molar-l rng.-*/n sec.-’/z in the pH range of 6 t o 9. ion, respectively, and Kn is the nth acidic dissociaThe reversibility of the electrode was further tion constant of the acid. tested by measuring the spontaneous potential of the dropping mercury electrode and of a quiet mer(7) 0 Schwarzenbach and H. Ackermann, Helo. C k i m . Acto, 80, 1798 (1917). cury pool in solution in which the mercury(I1) and
+
+
+
+
+
5784
JAMES
I. WATTERS,
JOHN
G. MASON '1ND 0 . E. SCHUPP, 111
the enta concentrations as well as the p H were varied independently over a t least a tenfold range. The two electrodes yielded essentially the same potential except for the iRcelldrop due t o the charging current of the dropping electrode. No drift in potential was observed with either electrode. Because the potentials of the quiet mercury pool versus the saturated calomel electrode were measured with greater precision, these values presented in Fig. 1 and Table I were used in the calculations.
I t was not necessary to consider HgzZ+even a t a p H of 2 since the addition of HC1 to the solutions resulted in no precipitation of Hg,C12. Substitution of the data into equation 5 yielded essentially a constant value of * 0 . 0 2 for Fl a t 25" and p = 0.1 in the pH range of 4 to S. This constant value of Fl is equal to Kloo. Typical data and results are summarized in Table I. Outside of the p H range of 4 to 8 the values of Fl began to deviate. TABLE I Curve 3 (solid circles), C,,,t;, = 1!1-.*211, CH,
A.
ESCE
$H
2 3e 2 91 :L2l 3.97 ,4.S!) 6.04 7.20 8.25 9.05 10.20 10.22 10 . 50 12,45
0.2
0.1
a,
$
( I . 23801 ,22883; 216113
,17471
. 12ti9fj -
-
0.0
-0.1
C.
2
4
6
8
1 0 1 2
@H Fig. 1.-Effect of varying the pH on the potentidl of the M mercury pool in the folloning solutions: (1) 1 X Hg(N03)2, 1 x 10-2 31 enta, 0 1 -11 E(XO3; (2) 1 X loe4Jf Hg(NOa)n,1 X 10+ 111enta, 0 1 dd KNOl; (3) (solid circles) 1X MHg(N08)2,1 X 31 enta, 0 1 M KNOa; (3) (open circles) 1 X M H g ( S O d ) ? ,1 X M enta, 0.1 M KNOa; (4) 5 X l o u 4.If Hg(S03)2,1 X enta, 0.1 M KNOa.
Comparing the potentials in curves 1 and 2, the potential shift due to a tenfold increase in mercury ion is +0.029 =t 0.002 v. in the $H range of 2 t o 10. Comparing curves 2 and 3 (open circles), the shift in potential due to a tenfold increase in enta concentration is close to 0.030 f 0.002 v. in the PH range from 2 to 12. In curve 3 the points for a tenfold dilution (open circles) lie close to those for the more concentrated solutions (solid circles) except i n the most alkaline range.
0
= log F s n
log I.']
06125 .(I1762 ,01357 .0454i . 10133 ,10296 ,11185 ,16371
B. Curve 2, Cent, = 2.60 0.21564 2.79 .20630 2.99 ,19240 3.90 145'28 4.74 ,09908 6.51 .05658 6.26 ,02040 7.1:3 - ,01320 8.70 - .Oil64 9.49 - ,10791 10. 1-4 - ,13325 - ,14152 I O . 54 - ,19085 12.52
I.
a v
4
- 0.2
VOl. 78
22.0:3 22.11 21 !):3 21.71 21. tih 21 .6i: 21.63 31 .6O 21 .no 22.88 22.92 23. 1!j 24.66
IIl
24.96 21. X7 21. 3:: 24.85
,
26.50 26.65 26,68 26.59 26.21
M, C H = ~
M
22.35 22.20 22.1s 21.80 21.65 21.64 21.63 21.68 22 .04 22.45 22.96 23. nz 24.53
21.85 24.85 25.02 25.19
27.12 26.88 26.80 26.46 26.01
Curve 3 (open circles), Celrta= 10-3 JI,GI, = 22,2!) 21.88 2.72 0.24266 22.01 24. is 3.00 ,22782 3.48 .19874 21.91 23.08 21.61 4.73 ,13156 5.35 ,09726 21 .BO 6.52 ,04082 21.61 7.99 - ,00601 21 .6(l 22.06 26.92 8.94 - ,04728 22.28 26.44 9.i3 - ,07452 22.62 26.42 10.15 - ,09287 12.57 - ,16064 2 3 . 53 25.98 F2 = F,, in acidic range and Fqb i n l x 4 e range.
-11
below a The constant value of Fza, 10:t',83* p H of 4, proves the presence of B single complex containing one hydrogen. This value of F2a corresponds to the value of the equilibrium constant K'llo shown in eq. 2 and I). Conversion to the equilibria in equation 8 yields IClio = 10i4,63 =t" , l o . Converting to Schwarzenbach's acidity expression, e ( ] .10 yieltls 1ia = I O 3 * . " ' . 'fhc
C0llst:lllt
VdlllC
of
F?,,,lO'!'i.fi')
'
'I,',.)
ill
tl1c
pH rang
