KOTES
,June, 1961 RATECONSTANTS FOR Ax At
THE
- = k(a Concn. of reactants, moles/i. Nickel acetate Pheophytin a
1 x 1 x 4 x lo-’
4 x 1x 1 x 10+
TABLE I FORMATIOS OF THE NICKELCnMpLEx OF PHEOPIIYTIN a CALCULATED FROM
- 2 ) 0 , 7( h - 2 ) O . 5
= k(a
TEMPERATURE RA~YGE i 4 TO 115’
OVER THE
Ratio constant
c
740
3 21 f 0 28 x 3 25 =k 20 x 3 70 =k 34 x l O +
900
1150
103O
1.93 & 0 15 x loe5 5 . 9 2 f 0 . 1 4 x 10-6 & . i o x 10-5 6 . 0 6 f .22 x 10-5 1.80 & .05 X 10-6 5 . 7 7 zt .13 X 10-5
1 78
for the four temperatures and the three initial concentrations of the two reactants. Extents of conversion were small, falling in the range of 2 to 25% for the reaction times involved, which ranged from only one to 32 hours so as to reduce the effects of side reactions which might have become important at higher conversions and to facilitate the summarization of the results by the use of initialrate approximations. To ascertain whether errors might arise through thermal instability of the complex, a series of measurements was made to evaluate the stability of the complex under the conditions of the rate measurements. In reaction periods twice or four times as large as the longest periods employed in the complexing experiments the recovery of complex in the degradation experiments was over 95%, except in the case of the ’74’test, in which the recovery was about 90%. Clearly, no significant degradation of the nickel complex of pheophytin a took place in the course of the complexing reactions. The coefficients in the general rate expression dt
1079
- z)a(b - s)P
were estimated by solving the two sets of simultaneous equations arising from the experimental data in which a or b (for nickel acetate and pheophytin a, respectively) was held constant while the initial concentration of the other reactant was varied ; in all cases dx/dt was replaced by the observed initial rate Ax/At. In this manner a series of evaluations for both 01 and P was obtained, the value for 01 ranging from 0.62 to 0.86 with standard deviations for each temperature ranging from 0.03 to 0.13, and the value for /3 from 0.46 to 0.66 with standard deviations from 0.04 to 0.12. The empirical ra,te constants for the complexing reactions at thle four temperatures are given in Table I. THE TEMPERATURE DEPENDENCE OF T H E ApD CORRECTION FOR T H E USE OF T H E GLASS ELECTRODE IN DzO
-
1 . 7 2 & 0.14 x l o - * 1.30 . i i x io-‘ 1.19 i .07 x
*
0.4 “pH” unit; i e . , pD = “meter pH reading” 0.4. More recent work by Glasoe and Long4 (who employ the symbol ApD for the numerical correction which must be added to the meter reading to obtain pD) and hfikkelson and Sielsenj have established the following facts concerning the value of ApD:
+
(1) The potential difference is due solely t o the glass electrode and not t o diffusion phenomena a t the liquid junction, etc.; (2) The value of ApD is not dependent on whether the electrode is pre-soaked in HzO pr. DsO; (3) ADD is not deDendent on the acidities of the solutions iivestigated 6ut remains essentially invariant from “pH 4-12”; and (4) The ApD value appears t o be independent of the type of glass electrode employed.
The present study to determine the temperature dependence (20 to 78’) of ApD was predicated on the great utility of the glass electrode in determining acidities and (as used in the pH-stat6) rate constants in DzO. The determination of these values is of particular importance in the study of general and specific acid- and base-catalyzed reactions. Experimental All measurements n-ere made with a Radiometer TTT-la pH meter employing a Beckman Calomel electrode and a Metrohm type X microglass electrode. The solution whose acidity was being determined was contained in a Metrohm type EA-662 microtitration vessel which was thermostated by water circulated from a Precision Scientific water-bath (zt0.1”). The cell compartment possessed three T openings which accommodated the glass electrode, the salt bridge and an ordinary T cap (without grease). To prevent rapid flow of KC1 from the salt bridge, a t higher temperatures, there was fused to the end of the salt bridge the wick tip of a Beckman Calomel electrode. Before the determination of acidity values the glass electrode was preequilibrated for 12 hr. a t the temperature to be employed. Standard solutions of 7.96 X S i 1.1% of DC1 in D20 ( repared from concentrated aqueous HC1 and !8.5y0 D$) and HC1 in H?O were prepared. The ApD correction was then taken as the difference in the pH meter reading for the DC1 and HC1 solutions a t each temperature. Determinations of ApD carricd out with solutions of 0.01 M DC1 in DzO and HCl in HzO were found to agree within 0.02 pH unit with the values of ApD determined a t lower concentration.
+
Results
Lumry,2 Hart3 and co-workers have reported that the glass electrode yields an apparent pD that is lower than actually exists in solution by about
In Fig. 1 are plotted the values of ApD us. l/’T. The linear plot fits the experimental points with an average deviation of =k0.008 pH unit and a maximum deviation of 0.02 pH unit. Included in the plot are the ApD values of Glasoe and Long4 determined a t 25’ with Beckman and modified Beckman electrodes and the value of Mikkelson and Nielson6 determined with a Radiometer glass
(1) To whom inquiries concerning the work ehould be addressed. (2) R . Lumry, E. L. Smith and R. R, Glantz, J . Am. Cham. SOC., 78, 4330 (1051). (3) R . G oHart, Nat, Ress Council, Cansda, CRE 423, June, 104%
(4) P. K. Glasoe and F. A. Long, J . P h y s . bhem., 64, 188 (1960). (5) K,Mikkelson and 6. 0. NielPen, z b z d . , 64, 832 (1960). (6) J. Pa Phillipa, ‘Autnmatin Titretorn,” Asadmmia P r m , New York, N,Y , , 1Od%
BY THOMAS H. FIFEAND THOMAS C. BRUICE~ Department of Chemastry, Cornel2 UnzLerszty, Ithaca, N e w York Heceabed December 14, 1960
1080
NOTES
Vol. 65
The second reduction wave of U(V1) was noted but it was not suitable for study since it merged with the hydrogen discharge ware and had poor definition. Using the Ilkovic equation, the diffusion current constant, I , was calculated. This “I” was not a true diffusion current constant as shown by the nonlinearity of the plot of i d us. C, but a t a given conI I I I I I J centration it is a convenient function to compare in 2.8 2.9 3.0 3.1 3.2 3.3 3.4 3.5 different media because it is corrected for changes I/OT, x 103. in ma/at1i6. Fig. 1.-Plot of the experimentally determined glass elecApplied potentials, as corrected for iR losses, trode correction (pH meter reading - p U ) in D20 us. 1/”T. Point 5! is the :orrection value determined by Mik- plotted 2’s. log [(id - i)/i]gave lines, the linearity kelson and Nielsen (22 ; ref. 5) and point 3 is that of Glasoe and slopes of which indicate the reversibility of the and Long (25”; ref. 4). reduction. Using Lingane’s’ equation, plots of El/, us. log electrode a t 122’. chloride concentration, gave straight lines parallel The formula which provides the electrode corto the abscissa in 0.001 M uranyl solutions in rection (=tO.O2 pH units) is water and the reduction apparently is reversible. I n various alcohol concentrations, there is first pD = pH. meter reading + 4.29 X 10’ - 1.04 (1) “T a slight decrease in I with chloride out to about ADDEDIN PROOF.-Extension of these studies has 0.02 M chloride, then a steady increase with inshown that equation (1) holds only for low resist- creasing chloride concentration. The viscosity ance electrocles designed for measurements a t am- of these solutions was found to increase slightly bient temperatures. Thus, in contrast to the with increasing chloride concentration so that it is Metrohm X (400 megohm a t 20’) electrode used seen the change in I cannot be due to viscosity here the ApD values for the Metrohm H (500 effects. This increase in I could be attributed megohm at 30”) electrode is 0.33 a t 78” and 0.29 either to the faster diffusion of the chlorouraniuma t 100”. For the latter determinations a specially (VI) complex compared to the aquo complex, or designed thermoststed deep cell with a vapor lock the faster disproportionation of the chlorourawas employed. This alteration in cell design did nium(V) complex, probably the former. Marked not alter the ApD values for the type X electrode. minima were observed for I in low chloride concenAcknowledgment.-This work was supported by trations which are hard to explain. The data for 21.6, 43.0, 79.3 and 92.2 weight grants from the National Institutes of Health and % ethanol in perchlorate media indicate the rethe National Science Foundation. duction to be reversible with one electron transferred. As the chloride concentration increases, POLAROGRAPHY I N WATER AND WA4TER- especially in ethanol-rich solvents, the reduction ETHANOL. I. URANIUM(V1) IN CHLORIDE appears to be irreversible and to approach a twoelectron reduction. This conclusion is drawn from AND PElRCHLORATE MEDIA I N ONE the slope of the El/, us. log [(id - i ) / i ]plots and MOLAR ACID the large increase of the diffusion currents. The plot of log [(id - i ) / i ]us. -E in 0.040 and BY m’IL1,IAM VES CHILDSAND EDWARD s. AMIS 0.050 M chloride in 92.2 weight % ’ ethanol gives a Chemistry Department of the L’ntvermty of Arkansas, Fayettevzlle, A-kansaa plot with two straight segments, with the slopes Recerwed December 28, 1960 of 0.044 and 0.063. This type of plot may be interpreted as meaning two reduction processes The polarographic reduction of U(V1) has been studied in 1 M acid solution in water, ethanol and with closely associated half-waves. Plots of - El/, ys. logarithm of the chloride water-ethanol media as a function of the chloride concentration, the ethanol concentration and the concentration, especially in ethanol-rich solvents, uranyl concentration. The polarograms were run show that El/, is strongly dependent upon chloride using a Sargent Model XV recording polarograph, concentration. The values of the ( p - q)/u potentials and cell resistances were measured factor in the Lingane equation’ are negative and, using a Leedis and Northrup Model K2 type poten- except perhaps for the limiting value in 92.2 weight tiometer and an Industrial Instruments Model yoethanol, are fractional. The values of -Ell2 a t selected chloride conRC16R a x . (conductivity bridge, respectively. centrations were found first to increase with inThe cell was designed to prevent contamination of the solution about the d.m.e. with material creasing weight per cent. ethanol out to about 25 from the salt bridge. KO maximum suppressor weight % ethanol, then to remain constant or decrease relatively sIowly depending on the molarity was used. of chloride, and finally decrease more precipitously Current rneasurements were made of the true current values, “imsX”’ a t the termination of drop depending on the molarity of the chloride (see Fig. 1). life by using the top of the recorder tracing. Beginning a t about 50 weight yo ethanol, I The diffusion current, id, was measured a t -0.4 v. decreases rapidly with weight per cent. ethanol in L’S. the saturated calomel electrode using a saturated lithium acetate solution as a salt bridge. (1) J . J . Lingsne, Chem. Recs., 291 ( l Y 4 1 ) .
