l I *ll WI "H W 14 266 THS U1 THE. SEHAWQE Ci? MCKELCH} $ALTS EN GLACEAL ACETEC ACED Thesis {or “no chmo 0‘ M. Sc EECEEQW STHE WEWESE'FY Jose-garb T. Lun$quist Jr. 1965 LIBRARY THE-.8“: Michigilli State fi Universit P33 \\\\\‘§§;“W‘ __ V V\ ‘5“ 0‘3“” TH E515 LIBRARY Michigan State University ABSTRACT THE BEHAVIOR or NICKEL(II) SALTS IN GLACIAL ACETIC ACID by Joseph T. Lundquist Jr. The preparation of Stoichiometric complexes of nickel(II) have been investigated. These complexes were NiClZ-ZHZO and NiBr2°2HZO. A complex from NiIZ of this stoichiometry could not be prepared. The complexes were prepared by diluting an aqueous solution of the salt with acetic acid until a precipitate formed which was isolated. Complexes of Ni(ClO4)2 and Ni(C2H3OZ)2 solvated with acetic acid were isolated and studied. Nickel p—toluenesulfonate was prepared in aqueous medium by re- action of nickel carbonate with a saturated solution of p—toluene— sulfonic acid. Drying the hexahydrate at 1300C. gave the anhydrous yellow salt. Near infrared and visible spectra and magnetic data suggested octahedral coordination for all complexes studied. Solubilities were determined for NiClZ, NiBrz, Ni(ClO4)2 and Ni(0Ts)2 in glacial acetic acid. The solubility of the salt in- creases with decreasing basicity of the anion except for perchlorate whose solubility fell between NiClZ and Ni(OTs)2. THE BEHAVIOR OF NICKEL(II) SALTS IN GLACIAL ACETIC ACID By Joseph T. Lundquist Jr. A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of MASTER OF SCIENCE Department of Chemistry 1965 VITA Name: Joseph Theodore Lundquist Jr. Born: May 2, l9bO in Bay City, Michigan Academic Career: Midland High School Midland, Michigan (l9Sh-19585. Michigan State University, East Lansing, Michigan (1958- ). Degree Held: B.S. in Chemistry, Michigan State University (1963). ACKNOWLEDGMENT The author wishes to express his gratitude to Dr. Kenneth G. Stone for his guidance and encouragement throughout the entire investigation and preparation of this thesis. ii TABLE OF CONTENTS Page INTRODUCTION . . . . . . . . . . . . . . . . . . . . . . . . 1 EXPERIMENTAL . Chemicals . . . . . . . . . . . . . . . . . . . . Standard Solutions . . Solubility Measurements . h_. . . . . . . . Magnetic Moment Measurements . Spectral Measurements O\O\O\U'LE"\.0LU Potentiometric Measurements . . . . . . . . . . . RESULTS AND DISCUSSION . . . . . . . . . . . . . . . . . \] Nickel(II) Chloride Dihydrate . . . . . . . . . . . . 7 Nickel(II) Bromide Dihydrate . . . . . . . . . . . . . l9 Nick€l(II) Iodide . . . . . . . . . . . . . . . . . . . 22 Nicke1(II) Acetate . . . . . . . . . . . . . . . . . . 23 Nickel(II) Perchlorate . . . . . . . . . . . . . . . . 28 Nickel(II) p-Toluenesulfonate . . . . . . . . . . . . . 33 SUMMARY AND CONCLUSIONS . . . . . . . . . . . . . . . . . . 3S LITERATURE CITED . . . . . . . . . . . . . . . . . . . . . . 38 LIST OF TABLES TABLE I. Analytical Data for NiC12‘2HZO . . . . . . . . . . II. NiCl2 Spectra (in acetic acid water solutions) . . . III. Solubility of NiCl2 in Acetic Acid . . . . . . . . . IV. Analytical Data for NiBr2’2HzO . . . . . . . . . . V. NiBr2 Spectrum . . . . . . . . . . . . . . . . VI.. Solubility of NiBr2 in Acetic Acid . . . . . . . VII. Spectra of Ni(C2HsOZ)2 . . . . . . . . . . . VIII. Analytical Data for Ni(C2H302)2'HC2H302 . . . . . . . IX. Analysis of Ni(ClO4)2°XHC2H302 . . . . . . . . . . . X. Solubility of Transition Metal Perchlorates . . . . . XI. Desolvation Rate Data of Ni(ClO4)2 8. SHC2H302 at 125°C.. . . . . . . . XII. Analytical Data for Ni(OTs)2 . . . . . . . . . . . . iv Page ll l2 19 2o 21 23 2A 29 29 30 3b LIST OF FIGURES FIGURE Page 1. Absorption Spectrum of Octahedral Nickel(II) . . . . lO 2. Solubility of NiClz and NiBrz in Acetic Acid . . . . 13 3. Thermogravimetric Analysis of NiClZ-2HZO and NiBrz‘Zfizo o o 0 o o o o o o o o o o o o o o o o o 15 A. IR Spectrum of NiClZ-2HZO (Nujol mull) . . . . . . . 16 S. The Structure of NiC12'2HZO . . . . . . . . . . . . . 18 6. The Structure of Cu(C2H3OZ)2-2HZO . . . . . . . . . . 26 7. Desolvation Rate Study of Ni(ClO4)2 8. SHC2H3OZ at 125°C. . . . . . . . . . . . . . . 31 INTRODUCTION There is very little research reported on nickel(II) salts in glacial acetic acid. Before any work can be done with nickel(II) salts in acetic acid one must know something about the physical prop— erties of these salts and the reactions these salts undergo in this solvent. There have been conflicting reports as to the stability of nickel acetate solutions in anhydrous glacial acetic acid. Campbell and Davidson (h) reported a solubility for nickel acetate of 12.37 mole percent and that above these concentrations a green white solid sep- arated out which was 5b—66 mole percent nickel acetate. The color of the solution was reported as green. Hardt and Eckle (12) reported that nickel acetate forms no stable solutions in acetic acid and that their solution was yellowish green which indicated the presence of acetonickolate. Tappmeyer and Davidson (32) found that the solubility of nickel acetate at 23°C. in glacial acetic acid was 0.557 mole per- cent. Above this concentration a green white solid precipitated out which appeared to be the hemi-solvate 2Ni(C2H302)2oHC2H302. Hardt and Eckle (12) found that the precipitation of nickel(II) ions with acetylchloride which yielded anhydrous NiClZ was nearly quantitative. However, about three percent [NiC14]2‘ was formed. It was also found qualitatively by Davidson (8) that nickel(II) sulfate was insoluble in acetic acid. This anhydrous sulfate was yellow and could be formed from the addition of concentrated sulfuric acid to hydrated nickel(II) nitrate, small amounts of water did not seem to interfere. Various solvated nickel(II) salts have been isolated from dif- ferent solvents. Hexasolvated nickel(II) perchlorates have been made with tetramethylene sulfoxide, dimethyl sulfoxide, and pyridine N— oxide by Meek, Drago, and Piper (23). Buffagni and Dunn (3) have iso— lated the nickel(II) hexadimethylformamide perchlorate and also nickel(II) chloride di-dimethylformamide. Wickandin and Krause (3b) have isolated nickel(II) perchlorate complexes with acetonitrile containing 6, h, and 2 acetonitrile molecules per nickel ion. These complexes were all octahedral with the perchlorate being monodentate in Ni(CH3CN)4(ClO4)2 and bidentate in Ni(CH3CN)2(C104)2. Other references for the ex- istence of bidentate perchlorate groups are given in this article (13,1h,25). Monnier (2h) isolated complexes of nickel(II) perchlorate with dioxane. The di- and monohydrates of nickel(II) chloride have been isolated by the stepwise azeotropic dehydration of the hexa— hydrate using ethanol as a solvent (19). Thermogravimetric data have also indicated the formation of these two complexes (5). Kolling (15) determined the dissociation constant for nickel(II) acetate in acetic acid. This was done by comparing the concentration and potential of a known base, sodium acetate, to the concentration and potential of nickel(II) acetate. Potentiometric measurements were made with a glass calomel electrode pair. He obtained a value of pK b for nickel acetate of 7.63. EXPERIMENTAL Chemicals The chemicals used in this investigation were generally reagent or A.C.S. Specification grades. Their names, commercial source and labeled purity are as follows: Acetic acid, Baker, A.C.S. Acetic anhydride, Matheson, Coleman, and Bell, A.C.S. Ammonium acetate, Baker Reagent Ammonium hydroxide, Mallinckrodt, 58% NH4OH, Reagent Cobalt(II) perchlorate, G.F. Smith, Reagent Crystal Violet, Eastman, Reagent Cupric perchlorate, G.F. Smith, Reagent Disodium ethylenediaminetetraacetate, Baker, Reagent Eriochrome Black T, Baker, Reagent Ferric perchlorate, G.F. Smith, Reagent Ferrous ammonium sulfate, Fisher, Reagent Ferrous perchlorate, G.F. Smith, Reagent Hydrochloric acid (37.3%) Baker, Reagent Hydriodic acid (h7%), Mallinckrodt, Reagent Nickel (ous) acetate, Baker, Reagent Nickel bromide, Amend, c.p. Nickel (ous) carbonate, Baker, Reagent Nickel (ous) chloride, Fisher, Reagent Nickel (ous) oxide, Baker, Reagent Nickel (ous) perchlorate, G.F. Smith, Reagent A l—B-octyl-h-benzyl—5-imino-tetrazolinium ditartrate monohydrate (15) p-Toluenesulfonic acid, Matheson, Coleman, and Bell, Practical Perchloric acid (70—72 percent), Baker, Reagent Phenolphthalein, Mallinckrodt, U.S.P. XII Phenosafranin, Eastman, C.P. Potassium acetate, Baker, Reagent Sodium acetate, Baker, Reagent Tetraphenylarsonium chloride, G.F. Smith, Reagent Zinc perchlorate, G.F. Smith, Reagent Zinc sulfate, Allied, A.C.S. specification. Standard Solutions Standard perchloric acid in acetic acid was prepare by diluting 8.5 ml. of 72 percent perchloric acid to about 1 liter with acetic acid and then adding 20 m1. of acetic anhydride with constant swirling. This solution was standardized after 2h hours by titration of weighed amounts of dried potassium acid phthalate. Crystal violet indicator was used to the blue—green endpoint, (three drops 0.1 percent solu- tion in acetic acid) (30). Standard potassium acetate in acetic acid was prepared by dis- solving 0.1 mole of dried potassium acetate in 900 m1. of acetic acid. The amount of water present in this solution was determined by Karl Fischer titration. It was removed by reaction with a stoichiometric amount of acetic anhydride. About four days were required for this reaction to proceed to completion. Again it was analyzed for water to make sure the reaction between acetic anhydride and water was complete. The solution was then diluted to exactly one liter. It S was standardized with the standard perchloric acid in acetic acid utilizing crystal violet indicator (31). Standard sodium acetate was prepared by dissolving 0.1 mole of sodium acetate trihydrate in 900 m1. of acetic acid. The procedure for the preparation of standard potassium acetate solution was fol- lowed from this point. Standard 0.01M disodium salt of ethylenediaminetetraacetic acid (EDTA) solution was prepared by dissolving 0.02 mole of the disodium salt in distilled water and diluting to exactly two liters. Standard 0.01M zinc sulfate solution was prepared by dissolving approximately 0.02 mole of zinc sulfate in 1500 m1. of distilled water. An aliquot of this solution was titrated with standard 0.0lM EDTA using Eriochrom black T indicator. The concentration of the zinc sulfate solution was determined and it was diluted with water until the concentration was exactly 0.0lM. Solubility Measurements Solubilities were determined by equilibrating an excess of salt with one hundred milliliters of glacial acetic acid. In certain cases acetic anhydride was added to remove water. Equilibrium was attained by heating the mixture for twenty-four hours at 700C. and then cooling at 25°C. for one week. The excess solid was removed by filtration. An aliquot of the filtrate was analyzed for water by Karl Fischer titration. Another aliquot was evaporated to dryness. The solid residue was dissolved in water and the nickel was determined complexo- metrically with standard 0.0lM EDTA solution, the excess being back 6 titrated with standard 0.0lM zinc sulfate solution utilizing Erio- chrom black T indicator (28). A solubility was then calculated from the amount of nickel present in the aliquot. Magnetic Moment Measurements Magnetic moments were determined using the Gouy method utilizing an Alpha Scientific Laboratories Incorporated electromagnet. Ferrous ammonium sulfate, magnetic moment 5.25 Bohrinagnetons, was used as a standard to calibrate the magnet and tube. All moments are at 2960K. Spectral Measurements Electronic Spectral data were obtained with a Cary (model lb) recording spectrophotometer. Infrared spectra of potassium bromide pellets and mulls of solid compounds were run with a Beckman IR5 recording spectrophotometer with sodium chloride optics. Potentiometric Measurements A Beckman Zeromatic pH meter equipped with a grounded solution shield, a calomel reference electrode and a glass electrode were used to make potentiometric measurements. The electrodes were stored in distilled water. Prior to use they were dried and rinsed several times with glacial acetic acid. No drift was noted upon standing in solu- tion so it was assumed the liquid junction potential remained constant. RESULTS AND DISCUSSION Nickel(II) Chloride Dihydrate The preparation of NiC12-2H20 from NiC12-6H20 has been accomp- lished by the addition of glacial acetic acid to an aqueous solution of nickel(II) chloride hexahydrate. The precipitation was carried out at room tmperature with stirring and a slow rate of addition of acetic acid. The most suitable concentration of nickel(II) chloride hexahydrate in the aqueous solution was found to be 1.0 gram per m1. of water. The amount of acetic acid required for a good yield of nickel(II) chloride dihydrate was 30.0 ml. per ml. of aqueous solu— tion. The precipitate was removed by filtration under atmospheric conditions and then dried at 105°C. for two hours to remove adsorbed acetic acid. This compound can also be prepared directly from the action of excess hydrochloric acid on nickel oxide in the presence of acetic acid and water mixture, followed by the addition of acetic anhydride. The reaction mixture contained: 5.0 grams of nickel oxide, 100 ml. of concentrated hydrochloric acid, 100 m1. of glacial acetic acid, and 50 ml. of water. The reaction was allowed to proceed with stirring until all the nickel oxide was converted to soluble nickel chloride. Acetic anhydride was added to this solution dropwise keeping the temperature less than 50°C. Upon the addition of 100 m1. of acetic anhydride a precipitate forms and by 200 ml. most of the nickel was precipitated as NiC12-2H20. The water is added to the above reaction mixture to accelerate the reaction of hydrochloric acid and nickel \] 8 oxide since the reaction in glacial acetic acid proceeds very slowly if at all. Each component of NiClZ°2H20 was determined analytically. Nickel was determined by a complexometric titration with standard EDTA solu— tion using Eriochrom black T indicator. The excess EDTA was back titrated with standard zinc sulfate solution. Chloride was determined volumetrically using standard silver nitrate solution with phenosafranine indicator. Water was determined by Karl Fischer titration. The sample of NiC12-2H20,however, was not soluble in the Karl Fischer reagent, but on standing for a few minutes all the water was apparently extracted and could be determined. The results are tabulated in Table I. Table 1. Analytical Data for NiClZ°H20 Sample % Ni % Cl %H20 1 35.Al A2.71 21.9 2 35.32 A2.81 21.7 3 35.32 A2.81 21.8 A 35.n3 A2.91 21.6 S 3S.A3 A2.82 --- 6 35.AO A2.78 --- 7 35.58 b2.SO --— Average 3S.h3 A2.77 21.7 Theoretical 35.80 h2.8b 21.76 9 The near infrared and visible spectra of nickel(II) complexes are quite indicative of the structure of these complexes. Jdrgenson (17,18) has determined the Spectra for a large number of octahedral nickel(II) complexes. There are generally four bands for these complexes in the ranges 8,000—11,000; l2,000—l3,000; 15,000- 19,000; 25,ooo-2s,ooo cm_l. The band at 8,ooo-11,ooo cm‘1 was as- signed as V1 and represents the transition from 3Azg(F) ——> 3T2g(F). This band therefore gives the value of lODq (27) which is indicative of the strength of the ligands surrounding the nickel(II) ion. Weaker ligands will split the 3d orbitals less than will stronger ligands which results in lower Dq values for weaker ligands. v2 represents the transition 3Azg(F) ———> 3Tlg(F). v3 is repre- sented by 3A29(F) ———> 1E and o4 by 3A2g(F) ———> 3Tlg(P). This 9 assignment was given by Brown (1) and Dunn (9) (Figure 1). A series of solutions with fixed nickel(II) chloride concentra- tion (0.031hM) and varying concentrations of water and acetic acid were prepared to examine spectroscopically the effects of acetic acid on the Spectrum of nickel(II) chloride. The solutions were run against reference solutions which were exactly the same minus the nickel(II) chloride. The data obtained are reported in Table II. Generally as the symmetry of a complex decreases the molar ab- sorbtivity, E , increases (6) and as the base strength of the ligands decreases the molar absorbtivity decreases. The data in Table II indicate that as the mole fraction of acetic acid increases the value of Dq,decreases and the molar absorbtivity increases up to a mole fraction of 0.82. At this concentration, however, NiC12-2H20 began to precipitate from solution. lO .AHHvHoxowz Hmncmzmpuo Mo ennpumam compaoomnd owpmmnmpomnmco AHI.EUV kocosvonm .H gunman coo.m ooo.oa ooo.ma 000.0m ooo.mm 1a w . _ 4 mm mm Hm.lll d. soda H) ma mm Adv Hm .niiu am ¢> O J KIIAIT qusqe JBION ll Table II. Nickel(II) Chloride Spectra (in acetic acid water solutions) M018 V1 V2 v3 V4 , Dq Fraction _1 _l _ _ E _ Acetic cm cm cm 1 cm 1 V4 cm 1 Acid 0.0000 8550 13800 15800 25250 5.20 855 0.23 8550 13800 15800 25250 5.30 855 0.56 8380 13600 15150 25000 5.65 838 0.77 8070 13150 18700 28200 7.00 807 0.82 7910 13000 18500 28000 7.50 791 0.99A0 7510 12500 13900 23200 1.AA 751 1.0000 7510 12500 13900 23200 1.81 751 NiC12'2H20 in a KBr -—— 12200 13300 22500 ——- —-- pellet As the waters of hydration are replaced by acetic acid or by bridged chloride the symmetry of the complex is decreased. symmetry is also decreased by distortion since these ligands are much weaker ligands than water. tance from the nickel ion than will water. Therefore, they will be at a longer dis— Since acetic acid is a weaker base than water one would expect that the molar absorbtivity would be decreased but the symmetry effects apparently over com- pensate this effect. At a mole fraction acetic acid of 0.9980 the molar absorptivity decreases to 1.88. hedral complex but shifted to lower energy. The spectrum is still representative of an octa— The Spectrum at a mole fraction acetic acid of 1.0000 is the same as that for mole fraction 12 0.9980. It can be assumed that the species in solution for these two solutions is identical. If one assumes the species in solution to be NiC12'8HC2H302, the symmetry of this complex would be only slightly less than that for hexaquonickel(II) ion, but the basicity of chloride and acetic acid ligands are much less than that for water ligands, Therefore, a lower molar absorptivity would result than that for the hexaquonickel(II) complex. A near infrared and visible spectrum of NiC12'2H20 was obtained by preparing a potassium bromide pellet of the solid NiC12-2H20. Only the wave length of absorption maxima could be determined from the Spectrum. These are listed in Table II. The spectrum appeared to be like the characteristic octahedral Spectrum for nickel(II)complexes. Solubility measurements of NiClz versus water concentration in glacial acetic acid were made. Table III contains these results. These data are plotted in Figure 2. Apparently the solubility of NiClZ is directly proportional to water concentration in glacial acetic acid. Table III. Solubility of Nickel Chloride in Acetic Acid Solubility H20 M/l M/l 0.0218 0.0000 0.0298 0.0392 0.0888 0.1050 13 cwo< owpmod cw cowpmppcoucou nopmz momno> mnmwz mo kpwflwbzaom ADV .Uwo< ompmod cw cowpmppcoocoo gowns mumnm> mfiowz mo xpwfifibsfiom Amv .m shaman A\z compmupcmocou booms 0H.0 m0.0 m0.0 N0.0 00.0 m0.0 40.0 m0.0 No.0 H0.0 00.0 _ 14. . fl 8 q .fil . m , . 00.0 i H0.0 01 O. 0 (A O. 0 *1/11 mummies no.0 mod 00.0 18 A thermogravimetric analysis of NiC12'2H20 was determined by the Dow Chemical Company analytical laboratories (Figure 3). The analysis was carried out in vacuum and the heating rate was approximately 2.5°C. per minute. The volatilization began quite slowly at approximately 96°C. The rate of volatilization increased until a rapid rate was attained. At about twelve percent weight loss the rate of volatilization decreased slightly indicating possibly a stable monohydrate. At 255°C. the weight loss was 22.3 percent which is larger than 21.76 percent water of hydration. This can be explained if one considers the volatilization of a small amount of hydrochloric acid leaving behind a basic nickel oxide. Also one cannot neglect the possibility of a small amount of adsorbed moisture. The infrared spectrum of NiClZ-2H20 Shows considerable hydrogen bonding, the hydroxyl peak at approximately 3.0 microns is split into two maxima, the lesser at 2.85 microns and the larger at 2.95 microns. The maxima at 2.85 corresponds to nonhydrogen bonded hydroxyl and the other at 2.95 corresponds to hydrogen bonded hydroxyl (Figure 8). A 3.5 gram sample of NiC12'2H20 was exposed to atmospheric moisture for seventy hours. A gain in weight of only 0.0080 grams was noted. However, when the humidity was very high the salt became hydrosc0pic. It was concluded that it was difficult for water to enter into the coordination sphere of this complex. The magnetic moment for nickel(II) in an octahedral field based on electron spin, angular momentum alone is given by the formula Magnetic moment = JFHTHI2) Bohr magnetons where n is the number of unpaired electrons. This formula gives a .Ommm.~nmwz mo mwmzfimcd ownpoew>mhm05nm£h AIIIV 15 Ommm.maowz mo mmmzfimcd omppoew>mcmoEnozH AIII¢ .m onsmmm .oo oompmuogeob 00m 04m OwH ONH o a 7 - . 0 boos onzm - 2% M“ 3 6 U. ''''''' 4" 'Ill 1. 7. O S N S pmoq o z m 4.0w Om l6 IOTDN IOTDN .Aaflse Honszv Ommm.maoaz go asobooam mH macho“: cm cpmcofim>m3 m a J W m .4 ouSmwm IOan TOFDN I l o 5% .a aoueqdimsuedl % 58 I C) N OOH 17 value of 2.83 Bohr magnetons for nickel(II). However, due to spin orbit coupling slightly larger values are obtained which apparently depends on the ligands surrounding the nickel(II) ion. A range of moments is therefore expected for octahedral nickel(II) complexes. This range is generally defined by the following limits, 2.9 to 3.3 Bohr magnetons (10). The magnetic moment for NiClZ-2H20 was found to be 2.98 Bohr magnetons, which is in agreement with the above range. This magnetic moment suggests that NiClZ-2H20 is an octahedral complex. The structure of NiC12-2H20 was determined by x-ray techniques (33) and is shown in Figure 5. " The above spectral and magnetic data are in agreement with an octahedral structure. l8 .oamm.~aoaz mo onsoosnom one .m unamau mm a: a: Ho 8 Ho 8 82 I \/ \ \ \ \ Hz \ / Hz \ / \ I x / .\ I .\ / .. x N aux N Iao\ o z o z Nickel(II) Bromide Dihydrate Preparation of NiBr2-2H20 from NiBrz-XHZO was done in a manner similar to that of the preparation of NiC12-2H20. One ml. of a saturated solution of NiBrz in water was diluted slowly with stirring with glacial acetic acid until thirty m1. of acetic acid had been added. During this procedure a light brown precipitate formed which was removed by filtration under atmospheric conditions. The precipitate was dried at 100°C. for three hours. This temperature should be held within five degrees for the volatilization of water begins at approximately 105-1100C. Analysis before drying showed less than seven percent acetic acid. Analytical data for NiBr2'2H20 are presented in Table IV. Ana— lytical methods parallel to those used for NiClZ‘2H20 were used to determine the various components in NiBr2°2H20. Table IV. Analytical Data for NiBr2'2H20 ,Sample % Ni % Br % H20 1 23.30 62.72 18.2 2 23.37 62.58 18.1 3 23.37 62.60 18.1 Average 23.35 62.63 18.1 Theoretical 23.10 62.65 18.17 The near infrared and visible spectrum of NiBrz in glacial acetic acid was similar to that of NiClz (Table V). The absorption maxima 19 20 are shifted to still lower energies. This would be expected since bromide is a weaker ligand than chloride (26). Table V. Nickel Bromide Spectrum. Mole Fraction 01 02 v3 v4 EV Dq Acetic Acid cm'l cm'l cm’l cm"1 4 cm‘1 0.9980 7350 12250 13300 22500 2.7 735 The infrared spectrum was also very similar to that of NiC12°2H20. Slightly weaker hydrogen bonding was noted in this compound. A thermogravimetric analysis was run on NiBr2'2H20 by the Dow Chemical Companyanalytical laboratories. The rate of heating was approximately 2.5°C. per minute and it was carried out under vacuum (Figure 3). The volatilization of NiBr2°2H20 began quite slowly at 70°C. This rate increased, held steady and then decreased until vol— atilization ceased. This occurred at 210°C. where the percent weight loss was 18.8 percent. This indicates either adsorbed moisture or volatilization of small amounts of HBr or both. This compound is less stable than NiC12‘2H20 as indicated by the different decomposi- tion temperatures. NiBr2'2H20 apparently does not have a stable monohydrate because no change in the rate of volatilization curve is noted. The solubility of NiBr2'2H20 was measured as a function of water concentration in glacial acetic acid (Table VI). These values when plotted as in the case for NiClz practically paralleled the curve obtained for NiC12(Figure 2). Rolling (20) reported in the case of 21 CdCl2 and CdBr2 and also in the case of PbClZ, PbBr2 and PbIZ that the solubility decreased with increasing base strength of the halides which is 01- > Br— > I- in glacial acetic acid. Table VI. Solubility of NiBrz in Acetic Acid. NiBr2 M./L. H20 M./L. 0.036 0.0000 0.0816 0.0820 0.0508 0.1050 The polarity of the nickel halogen bond decreases with electro- negativity of the halide. This makes displacement of the halide by acetic acid easier with increasing atomic number of the halide. The magnetic moment for NiBr2'2H20 was 3.28 Bohr Magnetons. This is in agreement with the expected range of moments for nickel(II) complexes in an octahedral field (16). The above data suggest that NiBr2-2H20 has a structure very similar to that of NiC12-2H20. Nickel(II) Iodide The preparation of NiIZ from nickel carbonate and hydriodic acid (87 percent) was accomplished by adding the acid to the carbonate which was in slight excess. Isolation was difficult because iodide is readily oxidized to iodine leaving behind a basic salt. A saturated solution of NiIZ in water was diluted with acetic acid which had previously had prepurified nitrogen passed through it to remove dissolved oxygen. No precipitate formed as in the case of NiCl2 and NiBrZ. But upon the addition of acetic anhydride a black precipitate formed which fits the description of anhydrous NiIz (29). Two methods were used to dry the material, oven drying at 90°C. and vacuum drying at room temperature. Both methods resulted in a product which was not pure NiIZ. The percent nickel was approximately 22.0 while the calculated value was 18.82 percent. It is felt that either nickel oxide or nickel acetate are present in substantial amounts. 22 Nickel(II) Acetate In an attempt to prepare a standard solution of Ni(C2H302)2 in glacial acetic acid 0.1 moles of Ni(C2H302)'8H20 was dissolved in one liter of glacial acetic acid. The solution that resulted was dark green. The visible and near infrared Spectrum of this solution gave the characteristic Spectrum for nickel(II) in an octahedral field as shown in Table VII. The position of v1 indicated that nickel was probably solvated to a greater extent by water rather than by acetic acid. The water concentration was 0.83 molar. Table VII. Spectra of Ni(CZH302)2. Mole Fraction -1 1 .. . cm cm cm cm D cm Acetic ACld V1 V2 V3 V4 q 0.975 8370 13600 15200 25000 837 1.000 8120 113150 18600 28800 812 0.0000 850 13800 15800 25250 855 To prepare anhydrous nickel acetate acetate solution, a stoicio- metric amount of acetic anhydride was added to the above solution to react with the water. After three days the water concentration was observed to be nill by Karl Fischer titration. The color of the solution was a green yellow. The visible and near infrared Spectrum of this solution still indicated octahedral coordination for nickel(II) but the absorption maxima in the spectrum were Shifted to lower energies as shown in Table VII. This would be expected if nickel were solvated by acetic acid rather than water since acetic acid is a weaker base than water. 23 28 It was observed that when thezmlydrous solution stood for a few days a light green fluffy precipitate formed. This precipitate approx- imated the composition of nickel acetate monoacetic acid. Similar results were obtained by Chappel and Davidson (8). Two methods were attempted to obtain a stoichiometric material, vacuum drying at room temperature, and oven drying at 110°C. Both drying methods gave a product which had slightly less than one acetic acid per each nickel. The analysis of each component was done by direct determination and the results are Shown in Table VIII. Nickel was determined as before by complexometric titration. Acetate was determined potentiometrically by titration in acetic acid as a solvent with standard 0.1M perchloric acid. Small amounts of water present in the acetic acid solvent did not interfere. Acetic acid was deter- mined by dissolving the sample in water and titrating with 0.05 molar sodium hydroxide utilizing phenolphthalein indicator. Table VIII. Analytical Data for Ni(Cz H302)2'HC2H302 % Ni % c211302 % 110211302 Calculated 28.81 89.85 25.38 Undried 28.52 89.35 26.2 Vacuum dried 25.80 51.86 22.2 Oven dried 26.80 52.97 20.5 All values are the average of three. The magnetic moment for this compound was 2.80 Bohr magnetons. 2S It is believed that this compound is a dimer similar in structure to that of the cupric acetate monohydrate (7,11) illustrated in Figure 6 because the magnetic moment is Slightly less than the value of 3.20 Bohr magnetons for nickel(II) acetate tetrahydrate. This low value is probably the result of partial 3d orbital overlap. A method for determining the dissociation constants for bases in glacial acetic acid (22) involved the use of a glass calomel electrode pair which have been shown to reSpond to the hydrogen ion activity in this solvent (16). A plot of potential versus the loga- rithm of concentration of the base gave for most divalent transition metal acetates including nickel acetate a slope of 0.5 which cor- reSponded to %§ millivolts per ten fold change in concentration (2). This slope suggests that the overall mechanism for dissociation of these bases is: > M(CZH302)+ + CZH302- Therefore one can use the following equation (2) to describe this mechanism: K + .12.... [H 1 = Vr___—E- Kb b Ks = autoprotolysis constant of solvent basic dissociation constant Kb Cb = total analytical concentration of base. Now if one uses a reference base and divides [H+] . by relerence + ' O O O I [H Junknown the following equatlon results. + 1/ [H ]reference = Kb unknown (Cb)unknown 2 [H+] Kb reference (bereference unknown 26 .ONmm.~ANOnm~oVso go onsoosnom one .0 unamwm ON: 0 \\\\\\\\\ so emu- \\\\\\\\\\ III/11111 0-0nm o o 0 0 £0. / \ 90$ 50 I 27 It follows that pr unknown pr reference 1 1 1 C 2 ' 2 ’ 2 09 C + 2 109 ApH = b unknown b reference 01‘ - log C pr unknown : ZAPH + pr reference + log Cb unknown b reference. This assumes the liquid junction potentials are equal. If one uses reference bases of known concentration one can determine a pH scale in acetic acid. For this work standard 0.001 molar solutions of sodium acetate, pr 6.58 and potassium acetate pr 6.11 were used (22). They were prepared by diluting aliquots of the standard 0.1 molar solutions of these bases 100 fold with glacial acetic acid. A small amount of water had to be present in the nickel acetate solution because the anhydrous solution is unstable. Apparently small amounts of water do not effect potentiometric measurement in glacial acetic acid but it may tend to drive the equilibrium. > Ni(CZH302)+ + CZHSOZ— Ni(CZH302)2 < to the right making pr appear smaller than it actually is. A value for pr of 7.63 has been reported and this work gave a value of 7.58 i 0.07. These values seem to be in good agreement. Nickel(II) Perchlorate A 0.01 mole portion of Ni(ClO4)2-6H20 was dissolved in 100 m1. of glacial acetic acid. The water concentration was 0.7 molar. To this solution 0.07 mole of acetic anhydride was added. The solution was mixed and it became quite warm (approximately fifty °C.) almost immediately because perchlorate ions accelerate the exothermic reac- tion between water and acetic anhydride. Within thirty minutes the reaction was complete and upon standing for about one hour a yellow green precipitate had formed. The color of the precipitate indicated solvation by acetic acid. The solution was allowed to stand for one week before filtration. Both the precipitate and filtrate were anhydrous as determined by Karl Fischer titration. The filtrate was also analyzed for nickel. The solubility of nickel perchlorate in anhydrous acetic acid was found to be 0.00185 i 0.00050 mole per liter at 23°C. Each component of the precipitate was determined analytically. Nickel and acetic acid were determined as before. Perchlorate ion was determined gravimetrically by precipitation with tetraphenyl— arsonium chloride. The data are presented in Table IX. Because a precipitate formed with nickel the above method was tried with cupric, cobaltous, zinc, ferrous and ferric perchlorates. With the above mentioned 0.1M divalent transition metal perchlorate solutions a precipitate formed, with the empirical formula M(ClO4)2-6-8 HCZH302.A111M ferric perchlorate solution on the other hand gave no precipitate. It is believed that these complexes would 28 29 Table IX. Analysis of Ni(ClO4)2-XHC2H302 % Ni % Clo4 % chhsoz 7.51 25.62 67.0 7.58 25.80 66.5 7.70 26.25 66.0 The above data suggested an empirical formula of Ni(ClO4)2(8-9)HC2H302. probably be good intermediates for the preparation of other complexes since acetic acid is a very weak ligand which should be easily re- placed. Solubilities of these transition metal perchlorates are listed in Table X. Table X. Solubility of Transition Metal Perchlorates in Acetic Acid Compound Solubility M/L. Co(C104)2XH02H302 0.007 Zn(ClO4)2XHC2H302 0.008 Cu(C104)2XHCZH302 0.009 Fe(ClO4)2XHCZH302 0.010 Fe(C104)3XHC2H302 > 0.1- 0nly two measurements were made for each perchlorate therefore, only rough averages are reported. It would appear from the above data that ferrous perchlorate could be separated from ferric perchlorate by this method. 30 For Ni(C104)2-8.5HCZH302 desolvation rate studies at 125°C. were made as shown in Table XI. The sample was placed in an oven at 125°C. and as time proceeded the percents nickel and acetic acid were measured as above and finally at the conclusion of the experiment each component was determined directly to see if decomposition had occurred. Table XI. Desolvation Rate Data of Ni(ClO4)2'8.5HC2H302 at 125°C. HC2H302 T1m€ Hours Ni 8.5 O 5.1 2.5 3.7 15.0 2.0 72.0 1.25 120.0 1.12 180.0 These data are plotted in Figure 7 and one can see that the final ratio of acetic acid to nickel approaches unity. The analysis of the final product indicates that no decomposition occurred. The percents of each component were; % Ni = 17.90, % 0104 = 61.10 and % HOAc = 20.58 which results in an empirical formula Ni(C104)2'1.12HC2H302. The material was liquid at 125°C. over the whole range of sol- vation. Only an increase in viscosity of the liquid was noted. It went from a light syrup initially to a thick paste finally. 31 OHN .oommH pm NOmmNUrm.w.NA¢OHUVHZ mo zuspm upmm cowpm>flomom mono: 08:. 02 QB 02 8 06 R .N. 6.8638 — . _ _ . d H O.w 32 In some recent work it has been shown that perchlorate was bidentate in some other low solvated complexes with organic sol- vents (38). If this were the case for nickel perchlorate solvated with acetic acid, which is likely since no shift in color is noted on desolvation, an increase in viscosity could be explained by bridg- ing of the perchlorate group forming polymeric species along with the fact that the "solution" is becoming more concentrated. Nickel(II),p-Toluenesulfonate To prepare nickel(II) p-toluenesulfonate a slurry of nickel carbonate in water was prepared. To this slurry, a saturated solu— tion of p-toluenesulfonic acid in water was added slowly with stir- ring. For simplicity HOTs represents p-toluenesulfonic acid and 0Ts_represents p-toluenesulfonate. Excess nickel carbonate was used since it could be removed by filtration leaving only soluble N1(0TS)2 in solution and a small amount of soluble carbonate. To minimize the amount of soluble carbonate as carbonic acid a vacuum was applied for thirty minutes. The dark green solution was then evaporated to dryness. An attempt was made to isolate a dihydrate of this salt by dis- solving it in water forming a saturated solution and diluting with glacial acetic acid as was done with the other salts. Five m1. of the saturated aqueous solution gave a precipitate which was light green upon the addition of 50 ml. of acetic acid. The composition of this precipitate approximated the empirical formula Ni(0TS)2’2H20. However, the water composition varied as the drying temperatures varied from 85 to 100°C. from more than two to less than two waters per nickel respectively. Drying at 130°C. gave a yellow compound which corresponded to anhydrous Ni(0Ts)2. The nickel was determined as before and the 0T3_ was determined gravimetrically by precipitation with 1-2—octy1— 8—benzyl-5—imino-tetrazolinium ditartrate monohydrate (15). These data are presented in Table XII. 33 38 Table XII. Ni(0Ts)2 Analytical Data % Ni % 0Ts 18.85 85.32 18.77 85.89 18.68 85.60 18.72 85.80 Average 18.76 85.85 Theoretical 18.65 85.35 Ni(0TS)2 was hydrosc0pic upon standing Open to the atmosphere. After one week the empirical formula for the compound was Ni(0TS)2‘ 6.60 H20 which was approaching Ni(0Ts)2‘7H20. The color of the Ni(0Ts)2'6H20 was a light blue but shifted over to a light green upon the addition of one more water. N1(0TS)2 seems to be very similar to NiSO4 (28). Anhydrous nickel sulphate is green yellow while the hexahydrate exists in two forms, a which is a blue form and B which is a green form. The heptahydrate is green. Solubility measurements indicate that Ni(0Ts)2 is quite insoluble in glacial acetic acid. Its solubility is 0.0016 i 0.0003 moles per liter at 25°C. with the water concentration of 0.0280 moles per liter. The magnetic moment was 3.20 Bohr magnetons. This indicated an octahedral structure. SUMMARY AND CONCLUSIONS Several compounds containing nickel(II) were isolated from aqueous solution by dilution with glacial acetic acid. They were NiClZ-2H20, NiBr2'2H20 and N1<0TS)2'XH20 where X is approximately two. NiC12'2H20 and NiBr2'2H20 were formed with the exact empirical formula stated. These halide complexes were found to have octahedral coordina- tion by SpectrOSCOpic and magnetic techniques. The infrared Spectrum Showed some hydrogen bonding in the solid state. Drying Ni(0Ts)2-XH20 at 130°C. gave the anhydrous salt. The physical properties of this compound were very similar to those of NiSO4. It was found with NiClz in solutions of acetic acid and water mixtures Spectroscopically that as the acetic acid concentration in- creased the band positions shifted to lower energies and the molar absorbtivity increased. These effects were explained through basicity of ligands and symmetry effects. Generally the weaker the base strength of the ligand the lower the molar absorbtivity will be. Also the band positions will be shifted to lower energy. The higher the symmetry is, the lower the molar absorbtivity will be. After a mole fraction of acetic acid of 0.82 (NiC12’2H20 begins to precipitate out) the molar absorbtivity decreases with increasing acetic acid concentration but the band positions still continue to shift to lower energies. The above effects can again be employed to explain this effect but now they oppose one another in explaining the low molar absorbtivity. 35 36 Apparently the base strength of the ligand is most important here. From anhydrous acetic acid Ni(C2H302)2'HCZH302 and Ni(C104)2' 8HCZH302 were isolated but the solvated acetic acids did not appear to exactly fit these empirical formulas. Ni(CZH302)2°HCZH302 probably has a structure similar to Cu(C2H302)°H20 which is a dimer because the magnetic moment is slightly less than what is expected for octahedral nickel(II). The pr 7.58 determined for nickel acetate in acetic acid was found to be in good agreement with another experimental value. Desolvation rate studies on Ni(ClO4)2°8HC2H302 showed that at 125°C. the non-solvated salt could not be obtained but that the empirical formula Ni(ClO4)2-HC2H302 was approached. This complex was probably octahedral and apparently contained bidentate perchlorate. Complexes similar to Ni(ClO4)é'8HC2H302 were prepared from the hydrated perchlorates of coba1t(II) cupric, ferrous and zinc. These compounds were quite insoluble in glacial acetic acid. A possible separation of ferrous and ferric perchlorates could be made in this manner since the solubility of ferric perchlorate was greater than 0.1 molar and the ferrous perchlorate solubility was less than 0.01 molar. The solubilities of the nickel(II) salts investigated here generally decrease with increasing basicity of the anion. It was found that the solubility increased in the order Ni(0Ts)2 < NiClz < NiBrZ. Per- chlorate ion should be the weakest base of all anions investigated therefore, it should have the largest solubility but its solubility fell between Ni(0Ts)2 and NiClZ. No solubility for nickel acetate 37 could be determined because no stable solutions could be prepared. This would probably be expected since acetate ion should be a strong base in this solvent. One would, therefore, expect a low solubility for nickel acetate. NiIz and Ni(0Ts)2 were prepared from the reaction of the cor- reSponding acid on a slurry of nickel carbonate in water. ExCess carbonate was used because it can be removed by filtration leaving behind a stoichiometric product. Pure nickel iodide could not be isolated but a product which con- tained some nickel oxide was isolated. Ni(0Ts)2 was isolated. Acetic acid would appear to be a good medium for preparing complexes since it is a weak ligand and could easily be displaced by slightly stronger ligands. The low solubilities of various salts should also aid in isolation of these complexes. Also anhydrous salts could probably be prepared by the reaction of the hydrated waters with acetic anhydride. These questions are unanswered and further work could be done in these areas. 10. ll. 12. 13. 18. 15. 16. 17. 18. LITERATURE CITED Brown, R. D., Quart. Revs., 6, 63 (1952). Bruckenstein, S., and Kolthoff, I. M., J. Am. Chem. Soc., 18, 2975 (1956). Buffagni, S. and Dunn, T. M., Nature, lgg, 937 (1960). Chappell, W. and Davidson, A. W., J. Am. Chem. Soc., 25, 3531 (1933). 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