The Oxidation Potential of Cobaltous-Cobaltic Salts in Nitric Acid Solution Citation Deahl, Thomas Jay (1936) The Oxidation Potential of Cobaltous-Cobaltic Salts in Nitric Acid Solution. Master's thesis, California Institute of Technology. doi:10.7907/BGPA-XC19. https://resolver.caltech.edu/CaltechETD:etd-04232008-095756 Abstract No abstract. Item Type: Thesis (Master's thesis) Subject Keywords: Chemistry Degree Grantor: California Institute of Technology Division: Chemistry and Chemical Engineering Major Option: Chemistry Thesis Availability: Public (worldwide access) Research Advisor(s): Noyes, Arthur A. Thesis Committee: Unknown, Unknown Defense Date: 1 January 1936 Record Number: CaltechETD:etd-04232008-095756 Persistent URL: https://resolver.caltech.edu/CaltechETD:etd-04232008-095756 DOI: 10.7907/BGPA-XC19 Default Usage Policy: No commercial reproduction, distribution, display or performance rights in this work are provided. ID Code: 1479 Collection: CaltechTHESIS Deposited By: Imported from ETD-db Deposited On: 25 Apr 2008 Last Modified: 03 Oct 2019 23:08 Thesis Files Preview PDF (Deahl_tj_1936.pdf) - Final Version See Usage Policy. 9MB Repository Staff Only: item control page

The Oxtdation Fotential of Cobaltous-Cobaltic Salts in Nitric Acid Solution Thesis by Thomas J. Deahl In Partial Fulfilment of the Requirements for The Degree of Master of Science in Chemistry The California Institute of Technology Pasadena 1956 The Oxidation Potential of Cobaltous-Cobaltie Salts in Nitrie Acid Solution The present determination of the cobaltous-cobal- tic oxidation potential represents a continuation of a series of investigations, besun in this leporatory under the direotion of Professor Noyes, on the oxidation poten- tials of strong oxidizing agents in nitric acid solution?’ **9 Previous work on this potential has been done by oberer, * by Jahn, ° and by Lamb and Larson. ® All these investiga- tors studied the same type of call: Pty H(1 atm.), H,80,(), /8,804(0) + ace \. Pt gf 804) g( 6p) Jahn states that Oberer found that mixtures of cobaltous and eobaltic sulfates in dilute sulfuric acid give a potential "far above 1.7 volte” azainst the hydro- gen electrode. In the cell studied by Jahn himself, the acid concentration (c) was 3 1 and the total cobalt con- centration (o,+ ¢,) 0.1 f. The potentials thus obtained were not corrected so as to refer the values to the stan- darg molal hydrogen electrode. However, Lamb and Larson have estimated the activity of the hydrogen ion in 3 n sulfuric acid to be 1.2, and, applying this to Jahn's measurements, they have computed a value for the single cobaltous-cobaltic potential from his data. In their own work Lamb and Larson used 4 n acid: they do not state the total cobalt concentration em~ ployed. Is order to obtain the single cobaltous-cobaltic potential from measurements of this cell, it is necessary to eliminate the potential of the hydrogen electrode in 4n sulfuric acid, This was done by measuring the elee- tromotive foree of a normal calomel cell against this hydrogen ealeotrode, a saturated solution of ammonium nitrate being employed as a bridge between the two half- cells. However, no corrections were made for the liquid- junction potentials of this cell. in order to avold excessive decomposition of the cobaltic salt, a11 measurements were made at 0°. How- ever, Lanb and Largon also made measurements at 16° and, by extrapolation, obtained a value for the potential at 25°. The work of Noyes and Garner®?® indicates, at iexst in the cage of ceric and thallic ions, that the tendency to form anion complexes ie reduced to a mini+ mam in nitrate solution; and accordingly, the potential obtained inf such solutions revresentea more nearly the molel oxidation potential. In view of theese facts and elso of the facet that in the vresent series of inves- 1,2,5 nitric acid solutions have been used, tigations it seemed desirable to study the cobaltous-cobsaltic potential in the same acid. Moreover, the present study is made under more varied conditions than those of previous workers. Therefore, there are desoribed in this paper the measurements at 0° and 25° which have been made on cells of the type co (e,) Cotte 5 Pt + ai(1 atm), HO10,(¢'), 1On( 6) + 1’ » Pt{or Au). Since the cobaltic salt decompoaes at a measurable rate at 0°, it hag been vossible to obtain scme information on the kinetics of its decomposition. hese results will be discussed briefly. Apparatus and Materials Preparation and Analysis of the Solutiona. - A stock solution approximately 0.5 ? in cobaltous nitrate was prepared by dissolving Merek's Reagent Co(NOg),°6 Hd (nickel free) in the proper quantity of distilled water. This solution was standardized gravimetrically (as CoS0,). It was found to be 0.6259 f. A stock solution of nitric acid was vrepared by diluting o.p. nitric acid (ep. er. 1.42). The resulting solution was standardized against a sodium hydroxide solution and was found to be 4.99 f. Stock solutions containing cobaltous nitrate and nitric acid in different proportions were prepared by weighing out those quantities of these latter two solu- tions and of distilled water calculated to give the concentrations desired. The sobalt in these stock solutions was par- tially converted into the trivalent form by electrolysis. The electrolytic cell used for this purpoge was equipped with a vletinum stirrer, which served also as the anode. The cathode, a small strip of platinum, was enclosed in & Vessel resembling a small Cooch funnel inverted. A sintered glass plate prevented diffusion between the anode and cathode compartments. A current of about 2? ampere at 6-8 volts was passed for a period 10-20% longer than that theoretically required to convert the cobalt entirely into the oxidized form. Since the electrolysis resulted in the transfer of some of the cobalt into the ¢sthode compartment, it was necessary to blow the cathode solution back into the anode com- partment after the current was turned off. Of course, since hydrogen ions were reduced at the cathode during the electrolysis, the seidity of the solution wes no longer that which it was made up to be. Accordingly, if the resulting solution were dasired to be, say, 0,1000 f in total cobalt and 5.0 f in nitric acid, the solution vas mede up initially 0.1000 f in cobalt and 3.1 f in acid. The solutions of perchloric acid for the hydro- gen half-cells were prepared by diluting the c.p. 604 acid, and they were standardized against e sodium hydroxide solution. Calomel was prepared electrolytically by the method employed by sllis.’ fank hydrogen was purified by passin it tirst through ea concentrated potassium hydroxide solution, then over solid potasciwa hydroxide, and Tinaliy over an electrically heated platinum wire. The Eyérogen Uali-Gell. ~ The hydro:en helf-cell used was Bimllar to that deseribed by Noyes and Garner.” Before entering the cell the hydrogen gas was passed through a class coil anc then it was bubbled through perchloric acid of the same concentration as that in the half-cell. The apparatve was immerasd in thermostats at 0.15°0.15° or 25.004C.05°. Method of Potential Measurements. - A Student's Type Leeds and Northrup potentioneter was used together with a high-sensitivity gaivanonater. After the completion of the eliestrulysise of e cobalt solution, the hydrogen half-cell was immersed in the solution. Previously the hydrogen electrode had been checked against a 1 n HCl calomel celi at the game tenperature. Potential measurements were made over a period varying from 25 to 120 hours, depending upon the rate of @ecomposition of the trivalent cobalt, this being the creater the more dilute the acid. At frequent intervals samplea were pipetted into previous- ly weighed flasks containing an excess of ferrous sul- fave. ‘The flasks were again weighed and the exeess ferrous sulfate titrated with permanganate. Since the pink color of the divalent cobalt interferec with the permanganate end-point, this was determined potentio- metrically, using a vacuum tube voltmeter. The concen- tration of trivalent cobalt thereby determined was sub- tracted from the known totel cobalt concentration in order to obtein that of the divalent form. In the case of the runs mete at 0°, the pivettes used in withdrawing the cobealt solution had been chilled previously, as were the flsaks containing the ferrous sulfate. Fotentiol measurements were made immediately preseding the ramoval of the samples and again immediately afterwards. For the O° experiments these checked to 0.1 millivolt. However, at 25°, where these frecuently differed by as much as 1.5 millivolt, a mean vas taken as corresponding to the potential at the time of withédrewal of the sample. After the completion of the run, the hydrogen half~cell was removed from the remaining cobalt solution and again checked against the calomel cell. the Observed end Caloulsatec Tlectromotive Forces Waoh of the following tables presents the data for one run. There are recorded tke observec potential measurements, the correcticns toa be epplied to these observed potentiais, and the formal electrode potentials reeulting therefrom. 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The gixth and seventh columms give the observed eotentials against platinwa end gold, respectively. ‘The activity coefficients (Wand &4') of the two acids are Given below their concentrations. ‘The values at 25° are those of Abel, Redlich, and v. Lengyei® vor nitris acid and those of Pearce and Helson® for perchloric acid. Inasmuch as there ara no data available at 6°, these same values are used at this latter temperature. The columns labelled 3 7" and ™ contain the Cc’ "2 H corrections which sre addea te the observed potential “obs to give the formal electrode yvotential 2°, The cuentity 4% = {Re/er) 1n( c,/¢,) is applied so as te provide for equiformal concentrations of co-~ bpaltous and cobaltic cobalt. ‘Since their activities are not known, the total concentrations Cy and Gp are uged. The quantity %, = (27,-1)(R2/F) infac/a'e!) L is applied in order to eliminate the licuid potential to the extent that this arises Pron the different ion activities ac and a'c' of the two acids. ‘The ion mobilities are here considered to be equal. The mean transference humber 4 of tne cation of the tyo scids nas been estimnted to be 0.65 at 2 and 6.87 at o%. 9 The quantity @, = (nT/F) lalate I/JBy, is sdded in order te refer the value to the standard molel hydro- gen electrode £, (1 atu), E’ (activity 1m). the formal electrode potential ®° thus obtainad is sconpidered to be equal to that of the cell + PttH, (1 atm), H (1 aj jimegta to 4 2)+r Pt{or Au}. | * Discussion of the fotentisl Leasurements The “MF in Relation to the Type of Slectrcode. - Jahn” experienced considerable difficulty in obtaining constant and reproducible results. Different platinun electrodes of the game area anc thickness which had been cleaned in the same manner, were found to give potentials differing by several centivolts. Moreover, the tine recuired to reach ecuilibrium with the golu- tion varied from a few hours to several days. nowever, he found that if on electrode were removed from a cobalt solution, the moisture schering to it allowed to evapo- vate, and the regidue remaining on the electrode ignitec te form cobalt oxide (pretskly Co,0,), reprodacibdle and fairly constant results could be obteined with it. Yiekel oxide behaved similarly. Lanb and narson® likevise used a platinum foil conted with cobalt onice. however, they found that « gold electrode gave a higher potential. iti thea present Investisatlon tae Sut 20 countered the same trouble as did Jahn. ‘whe divficulty lay lergely in the fact that electrodes were very slow in coming to equilibrium. However, it was found that the platinum stirrer, Which functioned es the anode in the electrolytic oxidation of the cobaltous solution, gave a reproducible and constant 7 if it were lightly platinized. Table I presents the data from a typical run using the platinum stirrer as the electrode. knowing that silver establishes its potential quite readily, Dr. U. b. Coryell suggested iis use as a potential mediator. Accordingly, a drop (1/28 m1) of In AgNO was added to about 225 ml of the oxidized cobalt solution. ‘The observed potential immeciately increased. A Second drop caused a further increase. After the addi- tion of about ten drovs, the potential no longer rose. The same effect was observed for a gold electrode, except that only about half as much AguO., was recuireac to attain the maximum potential. Figures 1 and & show graphically these effects. It will be noted that this maximum poten- tisl for gold is somewhat higher than that for platinum. In 1 n acid this difference at 0° is about 206 millivolts, while in 4 n acid it amounts to about 5 millivolts. Strangely enough, the difference at 25° in the latter acid is about 25 millivolts. As the trivalent cobalt decom- posed, this difference increased slightly. It is interesting to note that if not enough AgNo. were sdded to reach the maximum potential for plati- 1.800 ° Platinum {E90 Effect of adding AgNo., upon the observed potential 1.740 Go 0.10 f HNO, 3.0 f a8] Ln mn 5 6 7 8 9 10 il le 43 Pleatimum Effect of adding AgNO, upon the observed potential Co 6.075 f 1. 740 HNO, 2.0 ff } 2 3 4 5 6 F 8 9 LO Jl 12) 23 Drops of 1 n AgNO, added 22 num, but yet sufficient to attain thet for gold, the observed potential for platinum increasec over ao cartain period as the trivalent cobalt decomposed, whereas the observed potential for gold decreased in accordance with the Nernat equation. Before the addition of Agho, the potential ob- ferved for the gol¢ electrode wag usually, but not al- ways, higher then that for platinum. In the ease of 1 n nitric ecid, the gold electrode gave a ootential slightly lower then that for platinum. In 2 n acid the former was about 10 millivolts higher; in 3 n acid, about 25 milli- volts; and in 4 n acid, about 40 millivolta. In order to note the effect of the added Agno, upon 2, Table I should be compared with Tables IX and X. The amount of AgNO, solution which was added varied from 6 to 13 drops (1/28 m1) to e volume of about 225-250 mi of cobalt solution. The stronger the acid and the more rapidly the drove were added, the greater was the total number of drove required to reach the maximum potential. Rifect of Stirring upon the Observed Potential. - Before the addition of the silver nitrate, the observed potential with platineun was from one to two millivolts lower when the solution was stirred. After it was added, the potential with stirring was from three to five milli- volts higher than without, In the case of the gold eleo~ trode, stirring increased the potential half a millivolt, 23 both before and after the addition of the silver nitrate. Rifesct of Adding Nickel Nitrate. - It was thought thet nickel nitrate might have an effect similar to that of silver nitrate. Accordingly, one drop af Lu nickelous ritrete was added to a cobalt solution. The potential im- modistely dropped about 20 millivolts in the case of plat- jmum, while with e,lé the decrease wea only about 2 milli- volts. After the solution was allowed to stand ahout 20 minntes, thea votantials came back te their previous values. In view of the facts thet the gold electrode gives a slightly more consiant ©° then platinum, that stirring is practically without effect upon the potential, that it eones to equilibrium more repidly {ae indicated by the fact that Less sliver nitrate is recuired to attain the maximum potential), and finally, thet it gives a higher potential than platinun, the values recorded for gold are believed to represent more mearly the cobaltous~ cobaltic potential. Woreover, the staller difference fu *° in going from 1 n to 4 nu acid (0 millivolts for platinwa and 16 miliivolts for gold) favors the use of the values obtained with the .o1d electrode. the existence of this difference may be accounted fer in any of several ways. In the lower acid strengths there way oeeur o hydrolysia of the cobaitic salt by a reaction like Co'’’+ H,0 = coo’ + eg”. or with ine creasing ionic strength the activity coefficient of the ou oobaltic ion may inorease more rapidly than does that of the cobaltous ion. Or the cobaltous ion may form nitrate complexes more rapidly than does the cobaltiec ion with incressing nitrate concentration. However, since the complexes of the cobaltous ion are so unstable, 1° this last explanation does not seem to be a very probable one. It might be thought that a lerge part of this difference may be due to the inadequacy of the liquid junction correction @. Accordingly, in one run in v ad which 5 n nitric acid was employed, potential measure- ments were made first with a lu HCLO, hydrogen elec- trode and then with a 2n HOLO, electrode. The average of the H's caloulated from the potentials observed. with the latter electrode was only hal? a millivolt higher than those obtained using the 1 n electrode. Noyes and Yossiakoff~ have tested the reliability of the correction when the nitric and perchloric acids are both 1 molal (activity). In the following table ars summarized the re- sults of the present investigation, the values recorded being those for the gold electrode. Temperature O° O° O° O° o. 25° 25. HNO, 1.6 2.0 3.0 3.0 4.0 3.0 4.0 Total Co 0.05 0.075 0.05 O2 OL ell O.1 CoM foot — 0,36011.% 0.57=5.1 0.8604.7 0.52-5.4 0.31365 3.6022.8 3.9015.4 z 1.800 1.806 1.808 1.811 1.816 2.840 1.850 25 The values at 0° obtained by Lamb and Larson® from Jahn' 8° data in 5 n sulfuric acid and by themselves in 4 n sulfuric acid are 1.779 and 1.775 volts, respec- tively. At 25° Lamb and Largon obtained 1.817 volts. It is possible that the difference between these results and those of the present investigation may be accounted for on the basis of complex formation between the cobal-~ tic ion and the sulfate or hydrosulfate ion. However, due to uncertainties in eliminating the potential of the hydrogen electrode in 3 n and 4 n sulfuric acid, too much significance cannot be attached to this relatively small difference of 25 or 30 millivolts. The kinetic activity of the argentic ion appears to be considerably greater than thut of the cobaltic ion. When solutions containing approximately equivalent cuan- tities of cobaltic and manganous salts are mixed, there immediately appears a brown color in the solution, in- dicating the formation of manganese dioxide. After half an hour, the color of permanganate is apparent. However, if ea drop of silver nitrate solution is addec after the two solutions are mixed, the color of permanganate im- medictely appears. In the case of chromic and cobaltic salts, the color of the solution after 24 hours appears to have changed little, if any, from the greenish blue of the cobaltic ion. If a drop of silver nitrate is then added, the color soon becomes a yellowish green and after about 30 minutes the color is almost the pure yellow of 26 dichromate. Diseussion of the Kinetics of the Denomposition of Cobaltic Nitrate by Water Figures 5, 4, 5 and 6 present graphically the data of Tables VII, IX, X and XI. ‘The abscissae repre- sent the time in hours at which samples were withdrawn, the time of removal of the first one being taken as zero. The ordinates represent the reciprocals of the cobaltic concentrations and the logarithms of these quantities. It was found that, in general, the rate of decomposition is second order with respect to the cobaltic concentra~ tion during the first part of the run, and first order Guring the latter, this being the cuse whether or not Silver is present. Figures 4, 5 and 6 are typical of the results obtained. However, there are exceptions. In the case of Figure 3 the decomposition is first order throughout. In contrast the daia of Table VI give a good straight line when 1/(Co'') is plotted against time, indicating second order throuchout. In an attempt to find a function which fits the data, various ones were trieé. The following ecuations were investigated. ++ ++ +++ a(Co. ) = k, So + ky (Co"*") (1) a(Co Cottt Cottt aco) = Ky + ke on) (2) 27 +++ . & ee. = ky, (Comt+) + Ky (Cott) * (3) a(Gort* ++4- cot)" at av 2 Gott) rs 8 (Co™”) +++ +o+, 2 - ate) * k, (Go"*) + ¥. e (6) 0 With the limited data at hand, all were elin- inated excevt (5) and (6). ‘The cobaltic concentrations of Teble IX (Figure 4) were plotted against time. Then tangents were taken from the curve (to give the values of a( Cot?’ )/at at various cobealtie concentrations) and the values thus obtained divided by the corresponding cobaltic concentrations. These are the ordinates of the points of Figure 7. The upper curve corressonds to the ratio (cot )/( co" )* as abseissa, and the lower curve to the ratio (Cot'')/(Co*"). ‘The former appears to give a slightly better fit to a straight line. ‘These are equations (5) and (6) respectively. Table XI (Fg- ure 6) gives a similar greph, equation (5) also giving the better fit. The data of Table X (Figure 5) do not give a good Pit to any of these ecuations. It is by an analysis such as this that ecuations (1) to (4) were eliminatec. From such a graph EK, is obtained as the y-inter- eept and Ky as the slope of the line. 28 ol ( TU) ou 0S On of O¢ OT \ med +4 \ +44°9) 30T= qgueseid = fgnay 3. oc Sonn T smeig 06 OTT QO wy wie og*T OL*T 02° T 00° 0D) BoT- 29 (Ut) omTy of * cS 4 ey (a) 3 mc & SEN ~ | OP) Pp. q eet a eR VA a + Z0T- \ RIN OT G2 of Ge g 4429) /T OTL 02 *T an ORT OS*T 43°) 30T~ 30 oan ae (e) it 4% mi us ty * ra = C2 ap oO cama | ‘ 4 LN us ‘ 5 ee we * je de ade x uw ° at \ + eo Evsrey F ao Sera Sed ~ 5 a ae AF ~ iS 8) , Q ?° 7 v O & - us 4 Le ae -present ~~ 1/ (Co*4*) Veep 20) fT bene O 9) FOT- ho 1.60 1.40 1.30 80 Q ry O rot 31 32 T emiyg Saeed reer eres 870" 70 9 4°90) 7G 79) caecdny Oh og 02 Ot O 4409) 44.09) rm 0 ae < 2 ob as O8T en Te : o> oO 2 + ! fe) fo) : S: Q . Z ie : an eo . A net ae + ro) : (@) ZA 33 However, equations (5) and (6) are inadequate, for if they represent the true state of affairs, then both lines should have a common y-intercept. In order to explain the existence of the first order reaction, 1+ seems necessary to postulate a free hydroxyl radical. The following reactions would then exprega the mechanisn. tott® 4 on” = Cot + OF (slow) HO = 8° + of (rapid equilibriun) cot** + on = Cot + B® 4 40g (rapid) 200" «48,0 = ao + ont + do, The rate is then given by a(oe = k1 (Got**) (on) = my (Co***)/ (HT). Whether this is thermodynamically possible, it is difficult to say, since the free energy of formation of the hydroxyl radical in solution is not known. The second order reaction 1s explained by the following mechanism. Goo’? + HO = Got? + #0, (slow) acer 4 Hg coo*t + Got + ont (rapid eqhilibriun) Hal, + acott* = ont + 2cett + Q% (rapid) uee™* + of.0 «= Wot 4 but + 0, Assuming that the rate is simply proportional to (Coo**) ana substituting for it the equilibrium ex- ression od (o**} (x*) 34 one obtains the reaction rate expression te eet +++, 4 * (Go**) (n*)” It is to be noted thet there is an inverse hydrogen ion dependence in both of these rate expres- gions, which is in accord with experiment. It is to be emphasized that the data ure too limited to permit a satisfactory study of the kinetics of the decomposition. The main purpose of this inves- tigation has been the determimtion of the cobaltous-~ cobaltic potential. The reaction rate study has been incidental. However, it seems sufe to say that tne kinetics are rather complex and a more thorough study than the present one is necessary to permit any defin- ite conclusions to be drawn. In view of this fact further speculations re- garding the mechanism of the decomposition are omitted. Summary This paper describes measurements of the elec- tronotive force of the combination conprising as one half-call a solution of cobaltous and cobaitic nitrates in 1 to 4 n nitric acid and as the other half-cell 1 to 2 th perchloric acid kept saturated with hydrogen gas. The formal oxidation potential of the reaction Cotil + wiles coll nes been computed to have values at 0° Nal AH ranging from 1,801 to 1.816 volts with ineressing acid concentration. At 25° it has the value 1.842 in 3 n acid and 1.850 volts in 4 nacid. In a given eaeld concentration, the 2° has been Pound to remain constant within 2 millivolts when the ratio cot t/cottt is varied several fold. A preliminary investigation of the decomposi- tion rate of cobaltic salts in nitric acid solution in the presence of silver nitrate has been made. (1) (2) (3) (4) (5) (6) (7) (8) Nod On References Noyes and Kosslakoff; J. Am. Chew. Soe., 5%, 1246 (1985) Noyes and Garner: J. Am. Chem. Soc., (1936) - The Oxidation Potential of Verous-Ceric Salts in Nitrie Acid Solution. Noyes and Garner: J. Am. Chem. Soe., (1936) - She Oxidation Fotential of thallous-Thallic Salta in Nitric Acid Solution. Oberer: Dissertation, Foerster's Laboratory, Zurich (1903) éshn: 2. anorg. Chen., 60, 292 (1908) Lamb and Larson: J. Am. Chem. Soc., 42, 2024 (1920) ™Tllis: J. Am. Chem. Soe., 38, 747 (1916) Abel, Redlich, and v. Lengyel; 4. physik. Chen., 132, 204 (1926) Pearce and Nelson: J. Am. Chem. Soc., 55, 3060 (1955) Pauling: J. Au. Chem. Soe., 54, $88 (1932) Acknowledement The author wishes to express his thanks and appreciation to Dr. Noyes, under whose direction this work was carried out. He is also deeply indebted to Dr. C. D. Coryell for many helpful suggestions and eriticisms and for the general interest he has shown in the work.