The Equilibrium Between Cuprous and Cupric Ions and Copper Citation Nutt, Donovan (1918) The Equilibrium Between Cuprous and Cupric Ions and Copper. Bachelor's thesis, California Institute of Technology. doi:10.7907/nxx1-4431. https://resolver.caltech.edu/CaltechTHESIS:07232024-204159052 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: Unknown, Unknown Defense Date: 1 January 1918 Record Number: CaltechTHESIS:07232024-204159052 Persistent URL: https://resolver.caltech.edu/CaltechTHESIS:07232024-204159052 DOI: 10.7907/nxx1-4431 Default Usage Policy: No commercial reproduction, distribution, display or performance rights in this work are provided. ID Code: 16584 Collection: CaltechTHESIS Deposited By: Benjamin Perez Deposited On: 24 Jul 2024 18:14 Last Modified: 24 Jul 2024 18:14 Thesis Files PDF - Final Version See Usage Policy. 5MB Repository Staff Only: item control page

'.::he sis ~HE E(JJILI BTI IUM BETWEEN C.UPROUS Ji.HD CUPRIC IONS AlJD COl' l?ER BY Donavan Nutt Class of May 1918 DEP.ARTMEIIT OF CHEMISTRY ~HROOP COLLF.GE OF :'~(;FI!WLOGY, PASADSHA : ALU'. THE EQUILIBRIUH BETWEEN CUPROUS AWD CUPP.IC IONS A1'TD COPPER. The object of this inve :J ti gation is to determine the equilibrium constant of the reaction between cuprous and cupric salts and metallic copper, as represented by the equation ++ ~+Cu + = 2 Cu . With the aid of this constant, and the specific electrode potential of metallic copper against cupric ion, the specific electrode potential of metallic copp er against cuprous ion may be calculated by the relation E(cu+) = E(Cui"+) + RT log Ji. F Here E(cu+) denotes the specific electrode potential of cuprous ions, E( cu++) the specific electrode potential of cupric ions against metallic copper, and K the equilibrium constant of the above reaction. This equilibrium has already been investi gated by a nurt1be:r of chemists, Their results are not cone ordant, the values found varying from lJ. x 10-5 to 7 X 10-7. PREVIOUS WORK 7 Luther . . used a copper sulfate solution, and, after equilibrium had been established, determined the cuprous content by titrating with potassh.un 1.:;exmanganate. He found the constant, ( cu:) 2 , to have the value 7 x 10-7. Cu T 1 - Zeit. phys . Chem. 34, l+Sfs, (1900) and 36, 355, (1901). - 2 - .Abel 1 hung a strip of me tallic copper in cup ric sulfate s olutions of varying concentrations, which in addition contained about two equivalents of sulfuric acid per liter. The s.:rnount of cupr ous salt f ormed was derived from the de- crease in wei ght of t he copp er strip. He considered the questi ons of t he for mation of comp lex ions and of the valen ce of the cuprous salt. It was decided t hat experiments of this nature ccu1d n ot t hrow li ght upon the complexity of the ions, but the results indicate· that the cupric ion is Fo r the r a tio; (cuproµs salt)_2 , Abel obtained cupric s2.l t t he value, 4 x 10-5, at 25°, monovalent. ? Bodlander and Storbeck~ employed solutions of cuprous chloride in both hydrochloric acid and pot a ssium chloride solutions, in the attemp t t o determine t he complexity of the cuprous ion. Like Abel, they decided th at the problem coul d not be definitely settled by experi ments of this natur e. They a:ft er wards i nvesti gated soluti ons of cupr ous bromide. I n t hi s work the total copper was determi ned electrolytic al l y, and the cupric copper by titrating with s odium thiosulfate in the presence of a. l ar ge excess of potassi um i odide . cu:)ro us copper was then determined by difference. Th eir results confirm the conclusion of Abel) nam ely, that the cupr ous ion is monovalent. For t he ratio, rcu~ 2, they ~ obtained the value 5 x 10-5. l - Zeit. fur anorg . Chem. 26, 3s7, (1901) 2 - Zeit. fur anorg. Chem. 31, 1, and 400, (1902) The - 3 Chiu1 shook solutions of cupric perchlorate for several days with finely divided copper. One portion of thi$ equi- librium mixture was then used for the determination o:f cu;prou1:, ion by titrating with iodine, using carbon disilf:l.de as an indicator. A eecond porticin v,as taken for the electrolytic determlnation of total coppex, The cupric copper was deter- mined by difference. Considerable difficulty was encountered because of the fact.that the oxygen) dissolved in the very dilute iodine so1utions 1 tended to oxidize the cuprous ion to cupric. During the preliminary experiments it vrae found that the cuprous ion was oxidized as fast as it was r1.m into the iodine solution. Even after taking ~11 crdinary prscautioLe to overcome this source of error 1 the results obtained were not very concordant, varying from 1.1 x 10 -6 tc 1.9 x 10 -6 at 25°. Pre1lrninary experiments c-f the present investi gation indicated that equil:1.bri v.m is not eAtablished. even a;fte:r shaking the solutions for wore than three mcnthe i;;.t 25°. PR0CEDURE~1:IP]., OYED.:. In this investi gation the gener3.l outline of Chiu' s work has been followed, in that sclutions of cup ric perchlorate vrn re shaken with fine1y divided copper until equili bri:trn was established, and also that the total copper we.s deter- mined electrolytically. But a new me thod was employed for the deterr:,. ination of cuprous co:p;per. 1 - Mass . Inst. Tech. Theses, (1916). - 4 The method finally adopted makes use of the fact that cuprous ions reduce quantitatively r: J j_l ver nitrate to free silver in boiling ammoniacal solution. The pr oblem then we.,s an accurate determination of very small amounts cf silver. Two methods presented themselves. One 1 was the titration of silver solutions w'ith potassium iodide; in the presence of palladous nitrate, depending upon the fact tha.t an excess cf KI forms a precipitate of palladcus iodide~easily visible one part in a million. The objection to this method is the large volume of the solutions involved relative to the small concentration of the silver. The other method is likewise o.n iodimetric one, it being an adaptation of a. me:thod published by ?1!cLean a11.d Yan Slyke for the determination of small amounts cf halide as silver halide. The latter method was chosen f or this investigation. The method as developed by HcLee.n a."Yld Va.n Slyke 2 is as follows: To the solution of halide (or silver) to be titrated is added a solution containing sodi u.m citrate and sodium nitrate, with starch as an indicator. The silver 1 which had been dissolved in nitric acid was then titrated with standard potassium iodide. The first action taking place is the r:-,r ecipi tation of the silver as the iodide, according to the reaction As seen as there is the slightest excess cf potassium iodide solution the iodine is released by the action of the l Science 46, 1199, 622; (1917) 2 - Jour. 11\mer. Chem. Soc. 37, 112$, (1915) - 5 - sodi um nitr i te, and the s trong blue color, char acteri s tic of the iodo-starch r eacti on, is eas ily visible 1 2 KI + 2 FTI'J0 2 = 2 NO + 2 K OH + r 2 The purpose c:f the citrate is to regulate the acidity of the solution . About f our cc. of the reagent (se e below) are added to the silver solution before titration f or each gram cf free nitric acid cont a ined t herein. As four cc. of concentrated nitric acid were used t o dissolve the free silver, and l cc. of 16 W. acid cont ains app roximately l gram of free mm3, sixteen cc. cf t he reagent were necessary . The reagent itself was made up according to the forr.1uJ.a given by McLean and Von Slyke as follows: Trisodi tun Ci tr e.te, Na3C6H o 5H 2 o ....... . . 446. 0 grams 57 Sodium Nitrite, NaN0 2 ......... . .......... 19.0 Soluble Starch . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. 5 1later to .................. . ............. 1000 . 0 cc . It was f ound advisable in t his vrnrk t o employ a slightly stronger concentrati on of the starch than the f or mula calle d :for~ so to each mixtv.re j ust before titra tion five cc. of starch s olution were added. I n all other re sp ects the pro- ceedu.re a s given was follo wed. As may be at once apparent, the method developed has a mani:pulation very similar to the Volhard t i tration f or silver There a.re t wo very i mp ortant difference s however, interdependent upon each other. Fir s t, the reagent that is added and s econd, the Volhard titration requires one drop of a 0 . 1 N - 6 solution (about 0.05 cc.) for its end point, while because of the careful regulation of condl tions and the sim:plici ty of this method, a single drop of a O.Ol N solution gives an excellent and easily observed end point. It is at once parent that smaller vol~nes will allow of a sharper end point. To this end, the volumes of the solutions in the ti- trations were kept down to 50 co. at the beginning of the titration. The pipette employed for withdrawing the srunples was provided with a small tubular glass wool filter on the end. Preliminary experiments showed the necessity of this precaution in order to prevent fine metallic copper from being withdrawn with the sample. Any free copper on being boiled with the ammoniacal silver solution containing free silver, would tend, first, to oxidize the copper to cupric ion and, second, to displac:·e the equili bri ur.'l, with the greater formation of cuprous ion and so of free silver. Experiments with equilibrium solutions containing 1 milligram of metallic copper showed that a very material error ie introduced. by the presence of even this small amount of copp er. Another series of preliminary exp eriments were preformed upon solutions of arnmoniacal silver nitrate to determine whether the silver was quantitatively reduced from the same by the cuprous ion reaction. These were carried out by adding a measured quantity of a freshly standardized solution of cuprous chloride to a boiling solution of ammoniacal silver - 7nitrate. The silver precipitated out as free silver with that precipitated as the chloride was filtered off on a Gooch crucible. The free silver was dissolved in dilute nitric acid, and the filter well washed. remaining was weighed as such. The silver chloride The silver nitrate in the mixed solutions was preci pitated with hydrochloric acid and fil terecl off on the ec:.me G:occh filter. From the combined weights of the silver chloride, the excess of silver nitrate ws.s calculated. These results showed a quantitative reaction between the silver ion and the cuprous ion. The determination of total copper was in all cases made by depositing the copper out of a solution containing a definite concentration of nitric acid. This was found neces sary in order to obtain a clean pure deposit; depositions made from perchloric acid solutions were naturally inadvisable owing to the nature of the gas es eirolved, The proceedure followed was approximately t hat recommended by Talbot's Quantitative An alysis. PREP ARATIOlT OF THE SUB.§1ANCE S. Finely divided copper was made by electrolyzing a slightly acid~ approximately 0.3 N copper sulfate solution. ,..fi.. st.rip of copper served as the a node, and by employing a platinum wire as the cathode a high current density was obtained. The copper prepa.recl in this way can be removed by gently tapping the platinum wire. This copper was boiled - g - for a few minutes in a solution of coyfi er sulfate and sulfuric acid to remove traces of i mpurities that might have been deposited with t he coppe r, and also to remove any cupric oxide that mi ght have been formed by the action of the air, After boiling it was washed several times with dis- tilled water to remove the copper sulfate, Finally it was washed with 0.1 N cupric perchlorate, and then with a solution of the same concentration as that cf the equilibrium mixture into which it was to be placed. cf About ten grams the metal ao prepared was added to each ei.:;uilibri um bottle . Cupric perchlorate was chosen because of the fact that it is the salt of a strong oxygen acid, hydrolysis being thus less likely to occur. The cupric eal t .~was prepared by dissolving cupric oxide in pure perchloric acid. The cupric oxide was made by heating finely di v:i.ded copp er prepared as above in a Rose crucible. It was found that the action was never complete by heating in air 1 so the oxida tion was carried on in a current of cxygen. oxide wa s obtained, Thus a very pw:e copp er Thie cupric oxide was wei ghed out and a sufficient volUJi1 e of concentra ted perohloric acid added to cause the resulting solution to contain about ten pe rcent excecs acid. The presence of free perchloric acid serves the purpose of preventing the possible hydrolysis of the - -· salt. The standard stock s olution was made about 0.01 M. From the stock s olution the more dilute solutions were p re- pared. Then if hydrolysis does occur, the degree of hydrol- ysis is always the same. - 9 The containers employed for the equilibrium mixture of c1.1,1)ric pe rchlor e.te and copper were 250 cc. laboratory reagent bottles. Before use they were well washed out with a solution of cupric perchlorate of the strength to be contained in them. Before adding the copper and the cupric perchlorate solution, the bottles were filled with carbon dioxide gas. They were finally sealed with eealing wax, and placed in the thermostat. The thermostat was a metal basin of about 15 gallons ca.paci ty, provided vri th a wooden cover. The heating vni t was a 500 c.p. electric globe run continuously, and a 250 c.p. globe connected thru the mercury regulator. The regula- tor, after lengthy experimentation, proved itself capable cf maintaining the temp erature within one-tenth of a degree of 75° C or 50° C. The bottles were firmly attached to a central revolving shaft. METHOD OF PROCEDURE . . At the end of a predetermined fericd the bottles we re removed from the thermostat. Considerable difficulty was experienced during the preliminary experiments in finding an expeditious method of so securing and sealing the bottles that they might be safely opened before sufficient time had elapsed to change the equilibrium. Practice developed a. technique whereby less than li minutes intervened between stopping the motor and ta..'t(ing the sample from the bottle. ... 10 - As cuprous s alts are very easily oxidized by air, every possible precauti on to guard against this source of error was taken. The bottles were opened in an atmosphere of carbon dioxide. This gas was £reed from oxygen by pas- sing it thru two wash bottles containing cuprous chloride, hydrochloric acid, and :met allic copp er . The pipette employed to draw out the solutions was also connected with the carbon dioxide ge_n erator ~ well washed out and filled with the gas just previous to use. By an e.rrangement with a three way stop cock on the end of the pip ette, suction was a;pplied, and a sample taken il!J'i thout the solution coming into contact with the air. The sample was quickly transferred to a boil- ing soluti on of ammoniacal silver nitrate, care being ta.ken to keep the tip of the pipette under the surface of the hot solution. After boiling for three minutes to coagulate the silver formed by the reaction ~the solution was filtered thru a Gooch crucible, and the residue thoroughly washed first with boiling distilled water, then with a 50% solution of ammonium hydroxide, and fin ally with cold water. The silver on the filter consisted of very fine specks, often not visible to the eye. The condition of the metal enforced great care in the preparation of the Gooch filters to retain it. Only layers of minutely small fibers were successful for the purpos e. The silver so collected was dissolved in four cc. of concentrated nitric acid, in four portions, and the filter - 11 - washed free of silver nitrate. In case the volume of the solution, after the washings had been added, was over ~O cc. it was evaporated to that volume on the water bath. The cit-- rate reagent and the starch solution was added, s.nd the sample was ready for titration. The titrations were always carried out in a casserole . .Another sample was taken from the equilibrium bottle for the determination of total copper. Nitric acid was added to give the desired acid concentration, one gram of a.mrt1 011i um nitrate added, and the whole transferred to the platinum crucible. The time allowed fo1· the deposition was from nine to ten hours, and the current regulated accordingly. :CXPERI!£E1JTAL RESULTS. Experiments were run at both 50° and 75° C., as noted in the table below. By obtaining the equilibrium constant at these t~vo temperatures, the constant, and hence the specific electrode potential, at lower tempe ratures can be calculated. The equilibrium mixture was approached from both sides in both s eries of experiments. Checks were run on the determination of cuprous copp er during the first experiments with 0.0102 M and 0.00102 M solutions of potassium iodide. A later test of the stronger solution .showed the presence of Kro 3, and these results had to be discarded. not run on the total copper determination. Checks were - 12 - In correcting the value of the equilibrium constant for the dissociation of the salts involved, no data are available for the dissociation of the copp er salts of perchloric acid, so curves were plotted from the data given by Noyesl for the average value of the ionization of salts of the uniunivalent and unibivalent typ es. From these curves the values of the ionization of the various concentrations of cuprous and cupric salts herein involved were read off, and the specified correction made. The pipette used for drawing the sample for the cuprous copper determination delivered 99.3~ co. at 50° C., and that for the total copper sample 49.91 cc. In calculating the results the volumes of both were computed to 100 cc. other data wa s cal01.1.lated to mole p er liter. The The concentra- tion of the KI us ed in the titration of the experiments at 50° was 0.00102 M., and for those at 75° was 0.00106 M. 1 - Noyes and Sherrill - Gen . Prin. Chem.p. 56, (1916) ed. - 13 - The average values of the equilibrium constant of -the reaction ++ ... Ou+ Cu ~ 2 Cu. are 7.12 x 10~6 and 15.24 x 10-6 at 50° and 75° respectively. From these data may be calculated, with the aid of the van't Hoff equation, the heat of the reaction and the equilibrium constant at 25°. The heat evolved when the above reaction proceeds from left to right was found to be -6goo calories and . -6 . the equilibrium constant at 25° to be 2.9 x 10 The sp ecific electrode potenti al of cuprous ions against 25° may now be calculated by the aid of the equation given at the beginning of this report 1 . Employing -0.35 volt 2 copper at as the sp ecific electrode potential of the copp er, cupric ion electrode, the value -0.51 is obtained for that of the copper, cur)rous ion electrode. In addition, the specific electrode + ++ potential of the cuprous ion, cupric ion electrode, E(Cu, Cu ), may be calcule,ted by means of the relation E(Cu, Cu~)+ E(Cu+, Cu++)= 2 E(Ou, Cu++). The value thus obtained is -0.19 volt. 1- - Page 1. 2 - Lewis and Lacey. J. Jun. Chem. Soc. 36 , SO~-. (1g14) - 14 TABLE No . 1 No. Total copper 3. 4. .. 5. 6. 7. s. 9. 10 . 11. 12. 13. of gm .1100 KI 0.2475 0.3001 0.1716 0.1602 O.l4Sl 0.0960 o.16s5 0.2407 o.os94 o.054i 10.74 11.52 9.26 s.91 $,63 cc. 1. 2. c. c. .. 0.5995 0.3292 0.3173 • Z,up rous s al ts M./11 X- 10 5 9.61 9.02 10.84 6.64 ·• 5.12 17.10 12.11 12.17 1.102 1.1$2 0.9504 0.9065 0.8857 o.9s63 0.925g 1.113 o.6s15 0.5255 1.755 1.243 1.269 - - 50°. · Cupric salts M./11 Cu M./11 X- 10 5 X- 103 1.941 2.354 1.345 1.255 1.161 1.50s 1.321 l.887 1.083 1.15s 0.931 o. 6969 0.673 0.5196 1.703 1.21s 1.243 0, l.J.22$ 4.707 .. 2.5s2 2.489 o. $92 o.s73 0.96s 0.910 1.091 cu· -Xxl of M./11 x.10 3 1.4$1 2.0241.158 1.085 1.eo1 l.286 7.10 6.65 7.46 7.34 7. 50 7.29 1.133 7.16 1.590 6.s3 O.q5S 6.$7 0.3796 7.11 3.921 7.39 2.149 6.90 2.164- -7..:35 Mean .. 7.12 TABLE No , 2 - - 75° - 14. 15. i6. 17. l S. 0.2376 0.5261 0.3594 0.1785 0.1779 4.65 6. s7 h (6 .; • V 4. 09 3,99 .. 4. 959 7.327 6.237 4. 362 4. 316 1.844 4.104 4.849 2.794 1.3s2 5. g9g 7.144 . .. 1.37s 4.271 4.226 1.55g 15 .10 3.362 15. 1s 2.250 15.46 1.188 15.35 1.182 15.11 Mean 15. 24