The Molal Reduction Potential of a Selenium and the Free Energy of Selenious Acids From Equilibrium Measurements Citation Schott, Hermann Franz (1926) The Molal Reduction Potential of a Selenium and the Free Energy of Selenious Acids From Equilibrium Measurements. Bachelor's thesis, California Institute of Technology. doi:10.7907/6VJG-TD52. https://resolver.caltech.edu/CaltechETD:etd-11222004-144046 Abstract No abstract. Item Type: Thesis (Bachelor'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): Unknown, Unknown Thesis Committee: None, None Defense Date: 26 January 1926 Record Number: CaltechETD:etd-11222004-144046 Persistent URL: https://resolver.caltech.edu/CaltechETD:etd-11222004-144046 DOI: 10.7907/6VJG-TD52 Default Usage Policy: No commercial reproduction, distribution, display or performance rights in this work are provided. ID Code: 4635 Collection: CaltechTHESIS Deposited By: Imported from ETD-db Deposited On: 23 Nov 2004 Last Modified: 03 Oct 2019 23:37 Thesis Files Preview PDF (Schott_hf_1926.pdf) - Final Version See Usage Policy. 939kB Repository Staff Only: item control page

THE MOLAL REDUCTION POTENTIAL OF SEOENIUM ae ti AND THE FREY ENERGY OF SELENIOUS ACID FROM SQUILIBRIUM MeASURELENTS Thesis by Hermann F. Schott In partial fulfillment of the requirements for the degree of Bachelor of Science in Chemistry CALIFORNIA INSTITUTE OF TECHNOLOGY ~ Pasadena, California Purvose As there is among the free energy values recorded in literature no record of any values for the compounds of selenium, 1t wag thought advisable to determine tne free energy for one of the more important of these compounds. theory Thermodynamics leads to the following relation be« tween the molal electrode potential of a cell and the equillorium constant of the reaction taking place when the cell is in operation : ENF = RT ln K . Making use of tne fact that the potential & ip equal to the difference of the potentials £, and Eg , respectively , of the two nalf cells , E1 may be calculated , given Ee , and the equil~ ibrium constant K, at a definite temperature, T, accord- ing to the following equation in which the values of the various conss#ants have been substituted, and common log- arithims are used ;} iN = 0.5715 log K + HeN The molal free energy decrease is then obtainable from the following relation: “AP ,°= EallF And fron this the free energy of the oxidized compounds Can be calculated , if the free energy in the reduced . State is known . Choice of Reaction It was necessary to Tind a reaction in which selenium and selenious acid were brought into equilibrium by some suitable oxidizing agent, the molal reduction potential of which is already known . A further requirement was that no otner compound of selenium would be formed dur- ing the reaction . From some work done last year: by jir. Nordquist, in this labanatory, and fron the known properties of the element= , it was concluded that its reduction potential in acid solution would not differ greatly from that of ‘silver . The search for a suitable oxidizing agent in this neighborhood, which would not combine with selenium, re- sulted finally in the choice of iodine . Beckman and Faust, ( 2. anorg. Chem. 84 | 103) have shown that iodine and selenium form no compounds or solid solutions what- ever . | | The reaction of selenious acid with hydriodic acid has been studied for analytical purposes by Muthmann and sehifer , in 1825( Ber. 26, i008 ), and ag&in in 1695 by Gooch and Reynolds ( Z. fur anors. Chem 10,248), and fur- ther by Gooch and Pierce, 1696{ Z. anorg. Chem. di, 243 ). From this work it is evident that the reaction pro- ceeds to completion only in strongly acid solutions, if. at all, anda apparently does not take place in neutral solutions. The reversibilityof the reaction was read- ily demonstrated » As may be seen from the equation, se(s) + 3H-0(4) * 2 Ie{s) = HeSe0g + 4m + 4T- ~ rad - . the reaction is very sensitive to change in acidity or in iodine concentration, and therefore the equilibrium conditions can be adjusted, without great difficulty, to give analytically suitable quantities of the various con- stituents . Preliminary Experiments Last Spring several solutions of varying conpositions were prepared, and the equilibrium products were subject} ed to analysis . The analyses, for variaus reasons, were not satisfactory . However, sufficient data were avail- able for the approximate evaluation of an equilitrium constant in two’of the solutions, having widely different hydrogen ion concentration ( 1 N and .06N ) . About q0° "was found as the value for this constant for both of the solutions. The best analytical conditions as to conc~ ntration of all the constituents were those in the +06 N solutions . Analytical Methods In order to facilitate the analysis of such dilute solutions, it was decided to Search for a volumetric m method . The scheme of analysis finally chosen is based on the hydrazine method for selenium, which was Perfected by Strecker and Schartow ( Z. anal. Ch. 6& ,218) It seemed best to make up solutions with known iodide, selenious acid, and hydrogen ion concentration, so that the equilibriug value of the last could be indirectly det- ermined from the measured change in the other concentratvs ‘tions » without the necessity of an actual acidimetric titration in the dark colored equilibrium solution . It was found that the iodine could be titrated directly in the neutralized equilibriug mixture with either hydrazine sulphate, or sodium thiosulphate solutions , without any reduction of H,SeQ, provided that the titration as per- formed at room temperature and in a reasonable amount of time . An excess of hydrazine sulphate could then be add- ed to the solution , which after acidification would then be heated on the water bath until no mor e seleninmm ctre- cipitated. The solution would then be filtered, neutral- ized, and the excess hydrazine determined , from which the amount of selenium reduced could be calculated. The excess hydrazine could be determined by a nmber of dif- ferent methods. The results here were obtained by tite Yating the standard iodine solution in the presence of a buffer, according to Stollé ( J. pr. ch.(2) 9 ,33¢ (1g02) c-B (903 (2) 1340 ). The detailed procedure was as follows : Definite quantities of previously standardized sol- utions of HCl0s, KI, and HgSeOs, were introduced by means of burettes into a 250 cc. volumetric flask, and the sol- ution diluted to the mark with distilled water. The solution was then poured into a. bottle containing and ex- cess of solid phases of selenium and iodine , in an at- _ mosphere of carbon dioxide . The bottle was sealed , and shaken for several days at constant temperature ( 25 8) ° Two 30 cc. portions were then pipetted out ,. One pore ‘tion was then titrated with deci-normal thiosulphate sol- ution in the presence of a buffer . The sectind portion was treated with 50 cc. of hydrazine sulphate solution (.15M) sufficient buffer was added to cause ; disappear- ance of the iodine color, and the solution heated for several hours on the water bath . More buffer was then added , and the soluton filtered , and titrated with bodine , until the appearance of the yellow color . More reproducible results could be obtained in this kast titration without the use of the starch indicator . This is in accordance with the results of Kolthoff QJ. Amer.Ch. Soc. 46, 2009 (1924) ). ¥ " Both the hydrogen ion and the iodide ion concentrations were at first determined by noting the change in selenious acid content, and assuming that this all represented ant= oxidation or reduction , in accordance with the equilibrium reaction. - However, thehigh values for the lodide/tri-lodide raa&io, obtained in this way showed that this was ie not the case, but that some at-~ ‘mospheric oxidation had taken place, necessitating 4 different basis of calculation . As the solutions we re saturated with respect to iodine the values of the lodide/tri-Lodide-ratio should be nearly constant and equal to the values to be found in the literature for the same concentrations . In this way the todide con- centration could be determined from the experimentally found fodine concentration , and the change in hydrogen ton assumed to be equal to the change in fodide .This assumption 1s justified by the fact that the loss in {fodide from atmospheric oxidation, as well as in acc- ordance with the equilfibrium reaction is just equivalent to the loss in hydrogen ton . He0+Se+ Og * Hy Ses 4HI + O2 = 2Ie+ 2HeO Preparation of Materials O20 Two perchloric actd solutions were prepared by dil- uting the 60% C.P. acid . They were standardized against a@ sodium carbonate, preparé d by heating very pure sodium biearbonate ( Kah}baum's"zur Amalyse ") {n a platinum crucible at 300 C. for about half an hour, and drying over anhydrous calcium chloride . KI Merck C.P. KI was weighed out after powdering and drying at 200°C. , and cooling over calcium chloride . It was put into a volumetric flask, and ma@e up to — 250ce. with boiled water . Solution was caleutated to be .3749 N in KI on the basis of 99.88% purity. pS A pproximately deci-normal iodine solutions were made up and standardized against arsenious oxide . These were compared with the thiosulphate solution which was later used as a standard for other todine solutions. NRE, 30. A’ pproximately sixth normal was made up from some of the salt prepared last summer by Mrs Stone in this laboratory . The ratio of lodine to hydrazine was det- ermined and the strength of the hydrazine calculated from the known strength ef the todine . As a buffer -NaeHPQ: was chosen . This is preferable to NaHC0Os as it does not become basie upon boiling, and does not produce a gas, which might Amcrease the loss of iodine from the solution .« ‘HeSeOs Se "'Perten" , w ere dissolved tn concentrated nitrie acid , the solution evaporated to dryness , the residue redissolved in water, filtered, reevaporated, and sublimed onto a round bottomed flask filled with cold w& ter . A pproximately 4.5 grams of the white -meedles of SOethus obtained, wre then dissolved in la be) 2a. 2b 3a " P one half liter of water . Ten ec. portions were then ana lysed with the use of excess hydrazéne as previously described. te @ | The solid selenium was prepared by dissolving some Se0, ,prepared as described above , in water , adding HCL » and reducing with S 0, . The black amorphous pre- eipitate thus obtained was then filtered tn a Bhchner funrel, and washed with water and alcohol. I Commercially obtained, re-sublimed @.P. fodine was used . From the approximate eq utlibrium constant found last year the following sets of solutions were calculat- eds Final Concentrations Initial Céncentrations Se Ht Iv + I5 Sed, HY Iv + Is” 2010 064 2055 0.0 2023 2015 «020 064 «055 O12 O72 063 2013 2064 2052 0.0 e011 20 2013 2064 2052 2015 072 -060 It will be noted that each pair of solutions repres- ents the same equilibrium conditions approached from opp-. osite gidee of the reaction . They were a rranged in this way In order to prove q uantitatively the establishment of equilibrium . The actual initial compositions and those of the ey uilibrium are shown in the following tables: The concentrations are in mmols» per L. la. __ib 2a _ 2b 3a 5b *3576N HC10e cc. = ----- 50.00 ----- 50.00 ----- 50.00 . 1486N. HC104 « cc. 39.49 Steen 18.96 —“—— 20.18 ——— HC104 mmls/1. 23.29 71.53 11.27 71.55 12.00 71.53 “3749 KI ec. = 10,09 43.090 40.32 0 58423 KI mmls/l. 15.13 64.62 O 60.47 O 57.33 He SeOs mmls/1. O 112.237 90 14.90 3.98 17.94 Analysés of equilibrium samples: Volume in ce. 45.96 45.95 45.95 45.95 45.94 45.94 LO995N'NagSe03 cc. . 21.73 21.68 20.60 21.13 20.81 18.64 $Ie/mmls./l. 23.52 23.47 22.30 22.87 22.54 20.19 .[3- mmls.Z1. 22.19 22.14 20.97 21.54 21.21 19.86 +1523N NeaHgS04 ce. 50.00 50.00 50.00 $0%00 50.00 50.00 ,0992N La 33.08 . e1034N Lae 32.17 29-60 28.93 22.13 24.49 -Zoxid.ml.eq. 4.332 4.286 4.561 4.625 5.3526 5.082 ZtHeSeOs mmls/1. 11,70 11.59 15.68 13.72 17.72 17.59 Assuming no atmospheric oxidation: A-WaSé0s ml.eq./l. 4680 -2.60 54.32 -2625 54.91 ~1657 =I~ oom" 61.93 62.02 54.352 55.22 54.91 55.76 I~ sno 3974 39688 33635 36.68 35670 35.90 © Ht mo, 70.09 68.93 65.59 59.28 66.91 69.96 (I7)/(Te) 1.79 1.60 1.59 1.70 1.59 1.81 10 This table also shows the ‘odide/ tr'-indide ratios, ob- tained on the assimpticn that none of the oxidation was due to air. Calculations As one of the samples from each solution w as weigh- ed before analyzing, the concentrations could be convert- ed to molality - that is, from mols per liter to mols per 1000 grams of solvent . The activities of HI were found by using the activity coefficients found by Pearce and Fortsch (J- Amer. Chem. Soc. @x 45 2852 ) They give two sets of values, one based on the Lewis and Randall value of the mean activation of HCI st .005 nol » and the other on the Noyes and McInnes cvalue . For the sak e o8 consistency, the Lewis and Randall values were used , 4 as most of the other thermodynamic constants used were those of Lewis and Randall . The activation in each solu- tion was obtained by graphic interpolation of the tonic . strength from the table above referred to . This point is important , as the activation must be raisecd to the eighth power. The dissociation of selenious acid is calculated from the hyrgoen ion concentration and the constant g given by Blane . (J. de Ch. Phrs. 18 40) According to the werk of Blanc, K,* 2.7 x10" , K,=#5.1 x 10° . Ll The tri-lodide concentration is found by subtract- Ang from the analytically détermined iodine, the solubdil- ity of fodine . Lewis and Xandall givé the following table Total iodine salt concentration : .10m .02m .10m K (KI) /(KI,) : .99 1.04 1.06 Bray and McKay (J. Amer. Chem. Soc. 32 1207 (1910) ) have shown that the value of this constant does not de- pend upon the total todine concentiation, but only upon — that of the Lodide - “fhe The todide concentrations are then first approximate- ly determined from the tri-todide by assuming the value 1.02 for K in all se&utions . From these values the toti “4odide concentration ( that 1s, ZI=™+ IZ) , is read- ily determined for each solution , by intervolating these into the above table, K may be tore exactly determined , and the todine recaleulate’ from this value . }2 The following table gives the results of these calculations The concentrations in mmols per liter . io ib , 2a , 2 3a ig 22. 1s 22.84 20.97 21-54 21.21 £9.86 Zr Abe? 42.4 45.6 42.9 40.2 ar t aPF7RE-) i048 tléfg T2680 12680 12630 T.d22 KI “(in 1515 64.62 8 00.47 0. 57 +33 Rel04(4nit) 23629 71.53 11627 71.53 | 12.00 71453 H ( untor ) 58.96 5361 53.65 54.57 s4-04 Wt. solutes/1. 17-2 25-5 135696 22.6 14.7 21.9 ~ Wt. solution 1012. 1018. 1012. 1018. 1012. 10176 geaq./liter 994.8 994-0 997+5 995.0 997.5 994.8 Activation »859 2842 866 +342 0867 2843 HSe0s 66 68 075 «76 095 Peer) AgseOs(undis) 11-04 10.91 12693 12696 16677 16461 H 53-62 8952.29 54.38 556353 55-79 55.44 log (H) #2 7293 47184 07354 +7430 .7466 7438 log (I) +2 03543 = 6 3533) 63305) «63418 = 63351 3075 2 log « +1 8680 08506 .8750 .0506 8760 «3516 log fy a® +4. 9516 3225 23406 093554 3077 «9029 log(HeBeOs) $2 .0430 .0578 +1116 41126 $2245 .2204 9 colog W +8203 +0235 .0096 +0190 .0098 .0203 3 colog a Cle} 00030 =8.0015 °*9050 = +0011 #0050 log Kg 15¥ 08718 .7445 .885%3 .8768 1.0669 .b595 4 log CH) (T)o8413. 8064 6802 °7624 .7416 .8308 .6116 mean log K = -14.1166 K = %6 x 107*® aver. K = 7.65 x 10°™ E, + Ey = ~°2087 volts | Ey = .7444 volts % Subsequent Work Zn view of the high degree of dependence of these results on thr todide concentration, it was deemed ad- visable to determine the latter directly . For this-p “ptrpose the method of Bray and McKay (J. Amer. Ch. Soc 32 1193 (1910) ) seems to be well a da pted. The follewing s cheme of analysis w as tried 3; The sample is made neutral with a buffer, titrated to disappearance of the yellow color with hydrazine sul- phate. An excess of hydrazine sulphate is added , the solution slightly acidified, heated, and filtered as before. A small excess of potassium permanganate is “added, the solution allowed to stand for a minute or two, made neutral with a buffer, and the excess Lodine titrated with more hydrazine sulphate s-lution . The Solution, after evaporating to a suitable volume, is then acidified, titrated with permanganate, an‘ the iodine formed removed with carbon tetrachloride, and then determined with thiosulphate according to the ‘scheme of Bray an* MeclMcKay. | Some solutions were mad up, and analyzed by this” method - As yet, however, sufficient data for recal- evlation has not been collected. The results do show, howeverm that, due to the evaporation of todine, or seme other cause, as much as one or tro percent differ- ence may occur in the todine determination of tro supp- osedly identical samples. 14 This would account for she rather considerable variation of X in the previous results, since K 13 so largely de pendent upon the iodine concentration, as far as cal- culations are concerned. Conclusions Thidé work shows that the method of equittbrium measurements may be applied to such problems, and that fairly consistent results may be obtained, even with an inferior analytical method. This implies that the use of a more accurate procedure should yield very exact ve values for the electrode potential . Much aid was derived from the advise and suggestt - ions of Dr. Earnest H. Swift, under whom the work was started, and Yr. Don M, Yost , unde whose direction the major part of the work was carried on. Summary The equilibrium constant of the reaction was de- 3H, O (2) + 249+ Sess e0, §nI termfimed by direct analysis. A mean value of log K ws -14.1166, from which the electrode’ setSntial for sel enium of ~-T444 volts, based on the standard hydrogen eell,. The free energy of H, Se0,was found to be 100,975 calories ab 25 C.