Stoichiometric activity coefficients of sulphuric acid and the standard potentials of the lead dioxide/lead sulphate and mercury/mercurous sulphate electrodes at 25°C

Abstract

E.ni.f. measurements on cells containing aqueous sulpliuric acid at molalities less than 0.1 using hydrogen, lead dioxide/lead sulphate, and mercury/mnercurous sulphate electrodes have been used to derive stoichiometric activity coefficients of sulphuric acid. The new results and the previously reported measurements from 0-1 -8 m are consistent with activity coefficient ratios derived from vapour pressure measurements, but differ by 10 % from those reported in standard texts since these latter are based either directly or indirectly on the e.m.f. measurements of Hamer which are again shown to be incorrect. There is a systematic uncertainty in the determination of the activity coefficients due to uncertainty in the correct choice of value for the second ionization constant and to a lesser extent in the ion size parameter required for extrapolation to obtain the standard e.m.f.s.In previous communications 1 9 2 new measurements were reported for the cells :(1) (11) Pt, H2 I &SO4 (nz) I PbSO4 1 PbO2, Pt Pt, H2 I H2s04 (fn) I Hg2S04 1 Hg where m = 0.1 to 8 m. These two cells had been studied by Hamer 3 9 4 whose results are inconsistent with the vapour pressure data of Shankman and Gordon 5 and 6 with the isopiestic data of Stokes. By substituting values for the stoichiometric activities of sulyhuric acid solutions (derived through the Gibbs-Duhem equation from Stokes' vapour pressure measurements) into the Nernst equations corresponding to cells I and 11, it was shown,l, 2 that with the new measurements but not with Hamer's results, that constant values were obtained for the standard e.m.f.s of cells I and 11.Hamer 3 evaluated the standard e.m.f. of cell I by an extrapolation procedure but owing to the greater solubility of mercurous sulphate compared with lead sulphate this had not been possible 4 with cell 11. Harned and Hamer stated that below In = 0.05 the results for cell 11 became invalid because of the enhanced solubility of mercurous sulphate at low sulphate molalities giving rise to the possibility of an appreciable liquid-junction potential. In view of our interest 79 8 in other electrodes of the second kind where enhanced solubility at low molalities or other reasons may vitiate their use, new measurements were made on cell I1 to check Marned and Harner's statements which had not been supported by their experimental evidence. Parallel measurements at the same sulphuric acid molalities were made on cell I as a check and also on cell 111, since it was to be expected that, because of the higher solubility of silver sulphate compared with mercurous sulphate, cell I11 would fail at low sulphate molalities. 2050 Pt, H2 I H2S04 (m) I AgZS04 I Ag,

Thermodynamics of standard cells of the saturated cadmium sulfate type

“…This value is again large and has mistakenly lead some observers [66][67][68][69], under similar cases, to co nclude that the emf measurements are in error by as much as 1 mV whereas, in fact, the differences rest with the degree of resolution in the meas uring equipment.…”

“…'o Inspection also s how s that th e differences in I1G are practically negligible betwee n the fourth -and third-order fit s. If the preci- 10 Si nce th e va lues in tabl e 3 were ge ne ra te d fro m a thi rd-orde r equation rsee eqs (33) or (34)], the fit of a fourt h-ord e r e quation would be no bell e r than th at of a th ird-orde r. I-Ioweve r , de ri ved quantiti es . es pecia ll y dep , may be s~ghtl y diffe re nt ; act ua ll y the diffe re nce gives an es tim ate of the un certaint y in th e derived qu a ntit y. sion of the observations is lowe red to 0.01 m V (table 7) th e e rrors in I1G are still practi cally negligible between th e fourth -and third-orde r fit s (see footnote has mistake nly lead some observers [66][67][68][69], under similar cases, to co nclud e th a t th e e mf meas ure me nts are in error whereas, in fact, the differences arise from a lack of resolution or in imprecision of the instru· and a seco nd order fit to 0.01 m V in E (most common procedure reported in the literature); a difference of…”

Isopiestic determination of the osmotic and activity coefficients ofZnSO4(aq) atT= 298.15 K, and the standard potential of the electrochemical cell ZnHgx(two phase)| ZnSO4(aq)| PbSO4(s)| PbHgx(two phase)

“…As noted by Rard, (11) most modern studies and analyses yield values of K (HSO − 4 ) = 1.02 · 10 −2 to 1.06 · 10 −2 at T = 298.15 K, so the uncertainty from this source produces a maximum uncertainty of δ E o {cell (II)} = 0.4 mV. Clearly, these factors can account for only a small part of the E o {cell (II)} = 2.8 mV reported from the different evaluations at T = 298.15 K, (1,3,(8)(9)(10) and differences in the corrections for dissolution of Hg 2 SO 4 (cr) into the H 2 SO 4 (aq) must be responsible for most of the remainder.…”

“…This comparison indicates that although the derived values of E o {cell (II)} may vary considerably if the solubility and activity coefficient corrections are applied differently, the calculated E o {cell (II)} values are quite consistent when different sets of e.m.f.s are analyzed the same way. Similarly, Covington et al (8) studied the effect of varying the value of the hydrogen sulfate ion thermodynamic dissociation constant from K (HSO − 4 ) = 1.02 · 10 −2 to 1.08 · 10 −2 , which caused the derived values of E o {cell (II)} to change by E o {cell (II)} = 0.6 mV. As noted by Rard, (11) most modern studies and analyses yield values of K (HSO − 4 ) = 1.02 · 10 −2 to 1.06 · 10 −2 at T = 298.15 K, so the uncertainty from this source produces a maximum uncertainty of δ E o {cell (II)} = 0.4 mV.…”

“…A reinvestigation of cell (II) by Covington et al (8) yielded E o {cell (II)} = 0.6127 V at T = 298.15 K, again after conversion of their reported value to modern electrical units and to a standard pressure of 0.1 MPa. This value is 2.8 mV less than the value of Harned and Hamer.…”