A. L. BACARELA, E. GRUNWALD, H. P. MARSHALL AND E. L.
856
acid and for the hydrogen peroxide consumed. The symbol AEc"lM(materia1) denotes the energy of the idealized combustion reaction for 1 gram of the designated material. If the small concentrations of HzOz in the bomb sdutions are neglected, the values of AEco(manganese carbonyl) and of -2AE0[Mn(NO3)2. 10.309HzO] in Table I1 refer to equations I and I1 below. Two dilution corrections are necessary. The first, equation I11 below, corrects for the difference in concentration of the initial bomb solutions, and the second, equation I V below, refers each experiment to the same initial concentration of nitric acid, HN03(in 16Hz0).
+
+
Mnz(CO)lo 30.935HNO3~490.685H20 0.658Hz0 GO2 = 2Mt1(N0~)~.26.935HNO~. 493.342HzO f IOCOz AEoZ98.16 = -785.53 (I) 2Mn(N03)2~26.935HN03.493.342H20 = 2Mn(N03)2.20.618H20 26.935HN03.468.436HzO 4.288HzO AE02e8.i6 = -5.86 (11) 26.935HN03.468.436H20 = 26.935HN03.427.238H20 41.198H20 (111) AE0298.ia = 0.54 4HN03.GiHzO = 4HNOj.63.447HzO 0.553HzO (IV) AE'298.16 = 0.01
+
+
+
+
+
The sum of equations 1 through I V is
+
+
MnS(C0)lo 602 4HN03(in 16H20) = 2Mn(N03)2(in10.309Hz0) loco2 2H10
+
+
(VI
Vol. 62
PURLEE
TABLE I11 SUMMARY OF CALORIMETRIC EXPERIMENTS WITH MANGANESE CARBONYL A E ' (eq. I ) , kcal. mole-1
-785.53 -785.64 -785.32 -784.45 -787.17 -785.70
Dilution corn. - A E o (eq. 11), (eq. I11 & IV), kcal. mole-1 kcal. mole-1
5.86 5.96 5.50 5.84 5.44 5.64 AE'298.16 (eq. V), mean AH0298.16 (eq. V) Standard dev. of the mean Uncertainty interval
AEoae8.il (eq. V ) , kcal. mole-'
0.55 -779.12 .55 -779.13 . .55 -779.27 .55 -778.06 .55 -781.18 .55 -779.51 -779.38 kcal. mole-' -777.01 kcal. mole-' 0 . 4 kcal. mole-' 0.8 kcal. mole-'
from small thermochemical corrections, this value depends only on calorimetric experiments done in this Laboratory and the certificate value for the heat of combustion of benzoic acid used for calibration. Derived Results.-The values given in Circular 50@,1° or interpolated from values listed therein, for the standard heats of formation, in kcal. mole-', of HNOs(in 16Hz0), -49.18; Mn(NO&(in 10.309HzO), - 148.75; HzO(liq), -68.3174; and COdg), -94.0518, were used to obtain a derived value for the heat of formation of manganese carbonyl. lOC(graphite)
AED2se.ia = -779.12
+ 5oz(g) + ZMn(c, IV) = Mnz(CO)lo(c) AHf'2~8.10
= -400.9 kcal. mole-'
The results of the six pairs of combustion and Acknowledgment.-The Research Laboratories comparison experiments are given in Table 111. of the Ethyl Corporation have contributed financial The value of A H 0 2 ~ ~ .(equation le V) in Table I11 is support and a sample of purified manganese carthe direct result of the calorimetric studies; aside bonyl t o this investigation.
PROOF OF THE ACCURACY OF pH MEASUREMENTS WITH THE GLASS ELECTRODE I N THE SYSTEM METHANOLWATER' A. L. BACARELLA, E. GRUNWALD, H. P. MARSHALL AND E. LEE PURLEE Contribution from the Chemistry Department of Florida State University, Tallahassee, Florida Received April 9 , 1968
Previous claims of discrepancies between ~ K values A obtained potentiometrically with the glass electrode and conductometrically for acetic acid in msthsnol-water mixtures are shown to be incorrect. The two methods yield results that are identical with experimental error, thus establishing the accuracy of pH measurements with the glass electrod! in solvent compositions up to 95% (wt.) methanol. Experimental conditions which must be satisfied in order t o obtain accurate results with the glass electrode are reviewed.
Previous communications from this Laborat ~ r y have ~ - ~ emphasized that the glass electrode, when properly employed, correctly measures hydrogen ion activity in aqueous organic solvents. The glass electrode has been used in this Laboratory for the potentiometric determination of acid dissociation constants in methanol-water mixtures (including ~ K measurements A for acetic acid) and (1) This work was supported by a grant Foundation. (2) H. P. Marshall and E. Grunwald, (1953). (3) A. L. Bacarella. E. Grunwald, H. P. J. OTU. Chem., 20, 747 (1955). (4) E. L. Purlee and E. Grunwald, J. (1957).
from the National Science J. Chem. Phya., 21, 2143
Marshall and E. L. Purlee,
water-dioxane m i x t ~ r e s . 3 ~I n~ addition, we have employed the glass electrode in the potentiometric determination of ion-pair dissociation constants in the system dioxane-water.4-5 Direct evidence was reported by Purlee and Grunwald4 that the glass electrode gives a constant e.m.f. versus the hydrogen electrode in 70% ($.) dioxane-water over a wide range of hydrogen ion activity. Furthermore, comparison of e.m.f.'s measured with the glass electrode in 82% (wt). dioxane-water2 and in 70% (wt.) dioxane-water4 with e.m.f.'s previously obtained with the hydrogen electrode in these solutions indicated that iden-
A m . Chsm. Soc., 79, 1306 ( 5 ) E.
L. Purlee and E. Grunwald, ibid., 79, 1372 (1957).
July, 1958
ACCURACY OF pH MEASUREMENTS WITH GLASS ELECTRODES
tical measures of the hydrogen ion activity were obtained with both electrodes. However, in the syatem methanol-water, values have been reported recently by Shedlovsky and Kay6 for the acid dissociation constants of acetic acid which did not appear to agree with the results obtained in this Laboratory by potentiometric measurements with the glass e l e ~ t r o d e . ~Since these data were obtained by precise conductometric methods, the apparent discrepancies between the two sets of values have cast some doubt on the validity of measurements obtained with glass electrodes in media rich in organic components. In making comparisons with our results, Shedlovsky and Kay6 interpolated between their measurements which were made a t slightly different solvent compositions. Their interpolations indicated significant discrepancies with our results in the solvent composition range above 40% (vol.) methA anol. On plotting both sets of data ( ~ K versus wt. yo methanol) on large-scale, precision graph paper, we find no such significant discrepancies. Both sets of data fall on a single smooth curve with a mean deviation of 0.009 ~ K unit. A The best curve for our potentiometric data is identical, within experimental error, with the beet curve constructed for their conductometric data. I n Table I values obtained from the two best curves are compared with the actual experimental data. The mean value of the discrepancy, 0.017 p T i ~unit, is slightly less than the standard error in the potentiometric ~ K values.' A T o the best of our knowledge these results remove the last serious objection to the use of the (6) T. Shedlovsky and R. L. Kay, J . Phve. Chem., BO, 151 (1956). (7) We believe t h a t the apparent discrepancy between the two sets of data may have arisen from ambiguity in the definition of the term volume per cent. Shedlovsky and Kayo seem to have used volume per cent. in the manner described by Carr and Riddicksa
dMeoH1t where d is the density and t the temperature. In ref. 3 solvent oompositions were reported in terms of both weight per cent. and volume per cent. The latter was computed according to the definition
vel. % bIeOH a t t =
VMeOH
f
)
OH20
t
x
100
where v denotes the volume of the respect,ive pure component prior to mixing. In ref. 3, these density data were used a t 25': water, 0.99704, methanol, 0.78653.8b For any given methanol-water mixture, the two equations lead t o quite different values of the volume per cent. (8) (a) C. Carr and J. A. Riddick, Ind. Eng. Chem., 43, 692 (1951); (b) G. Scatchard and L. B. Ticknor, J . A m . Chem. Soc., 74, 3724 (1952).
857
TABLE I
COMPARISON OF POTENTIOMETRIC AND CONDUCTOMETRIC p K k VALUESFOR ACETIC ACID IN METHANOL-WATER MIXTURESAT 25.00' The values in parentheses were obtained by interpolation on large scale graphs. Wt. % MeOH
, Conduoto- P K A metric8
Potentiometric3
Difference
0.00 10.01 16.47 20.01 34.47 40.02 54.20 60.05 75.94 80.03 90.02 93.74 95,02
4.756 4.916 (5.025) 5.088 (5.367) 5.482 (5.798) 5.951 (6.521) 6.710 7.395 (7.881) 8.092
4.756 (4.910) 5.011 (5.071) 5.334 (5.458) 5,808 (5.954) 6.500 (6.703) (7.417) 7.858 (8.068)
0.000 ,006 ,014 .017 ,033 .024 .OlO .003 .021 ,007 .022 .023 .024
Mean 0.017
glass electrode in slightly aqueous organic solvents and, along with previously reported result^,^-^ confirm that the glass electrode can correctly measure hydrogen ion activity in these solvents. We have previously called attention to the experimental conditions which yield accurate results with the glass electrode, which are: (1) The glass electrode must be equilibrated with solvent of exactly the same composition as that in which measurements are to be made. (In 70% (wt.) dioxane-water we found equilibration to require ca. 24 hr.4) (2) When not in use the electrode must be stored in solvent of exactly the same composition as that in which measurements are being made. (3) The glass electrode must be standardized in solvent of exactly the same composition as that in which it is t o be used. (4) If measurements are to be made with cells involving liquid junctions, small but significant corrections must be made for the variation of the liquid junction potentials with ionic strength (see ref. 3, p. 758). For the most accurate work, the standardizing buffer must be chosen to resemble the components of the test solution as nearly as possible with respect to structure, charge type and concentration.
