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A kinetic study of the base hydrolysis of chloroamminebis (dimethylglyoximato) cobalt (III)

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A kinetic study of the base hydrolysis of chloroamminebis (dimethylglyoximato) cobalt (III)

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Content A KINETIC STUDY OF THE BASE HYDROLYSIS OF CHLORO- ' i AMMINEBIS(DIMETHYLGLYOXIMATO)COBALT(III) by Michael Fredericks )*) A Thesis Presented to the FACULTY OF THE GRADUATE SCHOOL UNIVERSITY OF SOUTHERN CALIFORNIA In Partial Fulfillment of the Requirements for the Degree MASTER OF SCIENCE (Chemistry) January 1968 UMI Number: EP41643 All rights reserved INFORMATION TO ALL USERS The quality of this reproduction is dependent upon the quality of the copy submitted. In the unlikely event that the author did not send a complete manuscript and there are missing pages, these will be noted. Also, if material had to be removed, a note will indicate the deletion. Dissertation Publishing UMI EP41643 Published by ProQuest LLC (2014). Copyright in the Dissertation held by the Author. Microform Edition © ProQuest LLC. All rights reserved. This work is protected against unauthorized copying under Title 17, United States Code ProQuest LLC. 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 4 8 1 0 6 -1 3 4 6 U N IVE R SITY O F S O U TH E R N C A LIFO R N IA T H E G R A D U A T E S C H O O L U N IV E R S IT Y P A R K L O S A N G E L E S , C A L IF O R N IA 9 0 0 0 7 C F This thesis, written by Michael Fredericks under the direction of h.i.$....Thesis Committee, and approved by all its members, has been pre sented to and accepted by the Dean of The Graduate School, in partial fulfillment of the requirements fo r the degree of MASTER OF SCIENCE Zpb/B* e. Dean D a te J/anujary^.,1968 .......... . THESIS COMMITTEE Chairman ACKNOWLEDGMENT I. wish, to thank my research, director,. Professor W. K. Wilmarth for his guidance and encouragement and Dr. K. R. Ashley for his invaluable assistance. ! CONTENTS I l I ACKNOWLEDGMENT ......................... LIST OF TABLES . . .................. LIST OF FIGURES....................... . Chapter I. INTRODUCTION II. EXPERIMENTAL III. RESULTS . IV. DISCUSSION APPENDICES I II. III. . REFERENCES iii LIST OF TABLES Table Page 1. Data for experimental run at hydroxide concentration of 4.164 x 10-1 M .................................... 16 2. The molar absorptivities of chloroamminebis(dimethyl- glyoximato)cobalt(III) at various wavelengths .......... 27 3. The molar absorptivities of the conjugate base of chloroamminebis(dimethylglyoximato)cobalt(III) and hydroxoamminebis(dimethylglyoximato)cobalt(III) at various wavelengths ..... 30 4. The molar absorptivities of aquoamminebis(dimethyl glyoximato) cobalt(III) chloride at various wavelengths............................................... 34 5. Data for the acid hydrolysis of Co (DH)2C1NH3..... 38 6. Data for the base hydrolysis of Co (DH)2C1NH3 ...... 40 7. Non-linear least squares analysis of data, all parameters free......................................... . 43 8. Non-linear least squares analysis of data. K fixed at the experimental value of 8.71 x> 10 M-1 ..... 43 9. Computer results for rate and equilibrium constants, all parameters free...................................... 56 10. Computer results for rate and equilibrium constants, K held at the experimental value of 8.71 x io1 Mt1. . 57 11. Calibration of the Vapor Pressure Osmometer with NaCl . . 63 12. Calibration of the Vapor Pressure Osmometer with sucrose...................... 64 LIST OF FIGURES jFigure Page i I 1. Sample potentiometric titration ........................... 10 | 2. Plot of -log(HA/A ) versus pH, compared with unit - slope, for acid-base titration of Co (DH)2C1NH3 .... 14 3. The plot of logfl-V/V^) versus time, for experimental values of Table 1 ........................................ 17 4. The molecular structure of chloroamminebis(dimethyl- glyoximato) cobalt (I II) . 20 5. NMR spectrum of the hydrogen bond of chloroammine bis (dimethylglyoximato) cobalt (111) ................. 21 6. NMR spectrum of the methyl protons of chloroammine- bis(dimethylglyoximato)cobalt(III) ................. 22 7. NMR spectrum of aquoamminebis(dimethyl- glyoximato)cobalt(III) chloride ........................ 23 8. Infrared spectrum of chloroamminebis(dimethyl- glyoximato)cobalt(III) .................................. 24 9. Infrared spectrum of chloroamminebis(dimethyl- glyoximato)cobalt(III) ....... .... 25 10. The molar absorptivities of chloroamminebis(dimethyl- glyoximato)cobalt(III) at various wavelengths .......... 29 11. The molar absorptivities of the conjugate base of chloroamminebis(dimethylglyoximato)cobalt(III) at various wavelengths . .......................... . . 32 12. The molar absorptivities of aquoamminebis(dimethyl- glyoximato)cobalt(III) chloride at various wavelengths' ...................................... 35 13. The molar absorptivities of hydroxoamminebis(dimethyl- glyoximato)cobalt(III) at various wavelengths .......... 37 v Figure Page 14. Graph of log of the observed rate constant versus pOH for base hydrolysis of chloroamminebis(dimethyl- glyoximato)cobalt(III) .................................. 41 I 15. The log of the observed rate constants and the calculated contributions of C and C, versus pOH ......... 54 cl D vi I. INTRODUCTION Many different complexes of cobalt(III) with dimethylglyoxime have been prepared. In addition to tris(dimethylglyoximato)cobalt(III), complexes have been prepared in which cobalt is coordinated to two dimethylglyoxime molecules and two additional ligands. These are complexes of the type [C0X2(DH)2]”, [CoX(DH)2A] and [Co(DH)2A2]*, where (DH) is the anion of dimethylglyoxime. X is a negative ligand and A a neutral ligand. All of the complexes with two dimethylglyoxime ligands appear to have trans configuration. Infrared and NMR studies that show the presence of intramolecular hydrogen bonding have been reported (1-3). This can only be present when the dimethylglyoxime ligands lie in a plane and therefore are trans to each other. There has also been a report of resolution of the complex Co(DH)2(0H)2 into cis and trans isomers (17). This report has not been confirmed. There is also a trans effect observed in this group of com plexes. A series of ligands can be arranged ranking their ability to facilitate substitution in the position trans to themselves. The order of reactivity of the ligands is 0H~ > NCS~ > N02” (4, 5). The compound chloroamminebis(dimethylglyoximato)cobalt(III) was prepared in 1906 (6) and since then the visible and ultraviolet q spectra of this compound have been reported (3, 7, 8, 9, 10). 1 2 Infrared studies which prove the trans configuration of this compound have also been reported (2, 3). No further studies have been carried out. The original objective of this work was to study the anation of aquoamminebis(dimethylglyoximato)cobalt(III) ion, and to continue work being done in this laboratory on other dimethylglyoxime complexes. The unsuccessful attempts to prepare this compound are explained in Appendix 1. One of the preparative methods tried was the base hydrolysis of the chloroammine complex. This was seen to be a fairly rapid reaction. The solid product obtained from this preparation had dif ferent properties the two times the reaction was carried out. It was then decided to study the kinetics of the base hydrolysis of the chloroammine complex. Base hydrolysis can generally be defined as an aquation reaction in which the product is a hydroxo complex. With the single exception of hydroxide ion, the major evidence shows no direct reac tion, in water, of a nucleophilic ligand with Co (III) complexes. Most examples of substitution reactions go through an aquo intermediate. Even very low concentrations of hydroxide ion greatly increase the rate of release of ligands from Co(III) complexes. In all cases in which a reaction of hydroxide ion occurs, the reaction has been found to be second order overall, first order in complex concentration and first order in hydroxide ion concentration. There are two main mechanisms which explain these results. The first is the S^ICB mechanism in which a proton is removed from an ammonia ligand forming a conjugate base. This conju gate base is an amido complex derived from the original ammonia complex, and is formed in a rapid acid-base equilibrium. This amido complex dissociates in a rapid but rate determining step to form a five coordinate intermediate. This intermediate adds a water molecule i! to form the hydro product. For the S^ICB mechanism there is the necessary condition that the molecule must have a moderately acidic j proton which can dissociate as described above. The behavior Of complexes lacking an acid proton is a strong argument for this ’ mechanism. For example, Co(CN)sBr~3 and Co(CN)sI-3 hydrolyse at a rate independent of hydroxide ion concentration over the alkaline ' range (11). The following are examples which strongly support the S^ICB mechanism. Firstly, in the reaction of trans-Co(en)2N02Ci2* with I - ! N02 , in DMSO, the rate is slow and independent of N02 concen- I I ' I | tration. Addition of small amounts of hydroxide ion causes the release of chloride ion to be greatly accelerated, the half-life changing j from 5 hours to 1 minute. The product in both reactions is trans-Co(en)2(N02)2+ (12). This result is explained by the S^ICB mechanism and not by the alternative S^2 mechanism, which can only give the hydroxo complex as the product. The second example uses the fact that the 180/160 ratio, in enriched water, is different in water and hydroxide ion. This then J labels the hydroxide ion in spite of rapid proton transfer. By j I measuring the oxygen isotope ratio in the product of the base j 4 hydrolysis it was seen that the hydroxo ligand must have come from a substituted water molecule, not hydroxide ion (13). Later work has been done on the base hydrolysis of Co (NH3)sX2+, where X = Cl", Br", I~, N03" in the presence of Y~ (21). Y~ equals N3~, OAc", N02~, NCS~ ions. This reaction 2 + gave a certain amount of the complex Cp(NH3)sY . This amount was essentially constant for each anion as required for a common inter mediate in the reactions. The common intermediate can only be ■ explained by an SN1CB mechanism. The base hydrolysis of trans-[Co(NH3)i , (x 5NH3)X]2+, where X = Cl", Br", N03" gives 50% of cis- and fran£-[Co(NH3)it (15NH3)OH]2+ (22). This result is consistent with the ls0/160 fractionation factors and the common competition ratios observed for anionic species and water molecules reacting with the intermediate. This also can only be explained by an S^ICB mechanism. The other proposed mechanism is the SN2 mechanism which involves direct attack by hydroxide ion. The greatest objection to this mechanism is that it is very unlikely that only hydroxide ion can act as a nucleophilic reagent with Co(III) complexes in water solution. Most of the experimental data can be fitted to both j. mechanisms. The previous examples cannot be accounted for by the S^2 mechanism. They are clearly explained by the alternative ^lCB mechanism. For this reason all reactions will be discussed in terms of the S..1CB mechanism. N The main aim of this project was to study the base hydrolysis of Co (DH)2C1NH3. It was expected that this complex would react j through an S^ICB mechanism since the complex has acidic protons on the ammonia ligand and on the dimethylglyoxime ligand. The complex is known to exist in a conjugate acid-base pair. The effects of this ’ dissociation will also be studied. II. EXPERIMENTAL Chemicals All inorganic chemicals used were of reagent, analytical, or primary standard grade as needed. The solution of sodium perchlorate was made by neutralizing the equivalent amount of sodium carbonate with 70% perchloric acid. The acid solution, pH 4, was boiled for one hour to expell dissolved carbon dioxide. The pH was adjusted to 10, and the solution filtered and neutralized. The concentration of the solution was determined by evaporation of a small aliquot to constant weight at 180°. The standard solution of silver nitrate was made by direct weighing of the analytical grade crystals (minimum 99.9% purity), which had been heated for two hours at 120° and cooled in a dessicator. The standard solutions of sodium hydroxide were prepared from carbonate free, 50% analytical grade solution. The solutions were titrated against primary standard grade potassium hydrogen phthalate, using phenolphthalein as the indicator. The standard solution of the disodium salt of ethylenediaminetetraacetic acid, (EDTA), was prepared by direct weighing of the analytical reagent after drying the crystals at 80°. This solution (0.015 M) was also titrated against 0.01 M zinc nitrate solution made from 99.99% zinc metal. The indicator used was 6 7 Eriochrome Black T. The concentrations from both determinations corresponded within 0.1%. Preparation of Compounds Chloropentaamminecobalt(III) dichloride was prepared by dissolving 50 g ammonium carbonate, and 25 g cobalt chloride in 250 ml water and 125 ml concentrated ammonium hydroxide in a liter breaker. A current of air was then blown through the solution for three hours. After adding 5 g ammonium chloride, the solution was boiled down to a syrup. Dilute hydrochloric acid was added to decompose the carbon ate, and the solution saturated with ammonia gas. After adding excess concentrated hydrochloric acid, the product precipitated out. The crystals were washed with a little dilute hydrochloric acid, 95% ethanol, and then dried in air at room temperature. Chloroamminebis(dimethylglyoximato)cobalt(III) was prepared by mixing 40 ml water, 5 g chloropentaamminecobalt(III) dichloride, 4 g dimethylglyoxime, 15 g ammonium acetate, and 1-2 ml concentrated acetic acid in a 1,liter beaker. The mixture was warmed to 50° with vigorous stirring until the dimethylglyoxime appeared to dissolve. At this stage the mixture was a red slurry. The preparation was always successful; however, the time of reaction varied from five to thirty minutes. The compound always remained insoluble, the color of the slurry changing from red to brown during the course of the reaction. Increasing the temperature did not increase the rate of the reaction. If the temperature was increased too much, the product was not brown but yellow. This yellow product was thought.to be 8 diamminebis(dimethylglyoximato)cobalt(III) cation. The chloroammine complex was filtered, washed with cold water and 50-60 ml absolute ethanol. The product was purified by dissolving as much of the complex as possible in water at 30°. This solution was filtered, frozen solid and allowed to thaw. The product precipitated out as large brown crystals, which were dried in air at room temperature. Analysis The compound chloroamminebis(dimethylglyoximato)cobalt(III) was analyzed for weight loss on heating, nitrogen, cobalt, and chloride. The weight loss on heating at 110®~120° was assumed to be water content. Spectral investigation showed no decomposition within experimental error at these temperatures. The weight loss was 2.7% and corresponded to % H20. An aqueous solution of the chloroammine complex was passed through cationic and anionic ion exchange columns in the Na+ and Cl" forms. They were Dowex 50W-X4 and Dowex 1-X8 respectively. No bands were observed from the compound since it is a neutral molecule". There fore the following analytical results are for a single compound. The compound was analyzed for nitrogen by the Kjeldahl-Gunning method. Approximately 0.15 g complex was decomposed by 50 ml fuming sulphuric acid with 10 g potassium sulphate to raise the boiling point, and 0.5 g cupric sulphate as a catalyst. The solution was boiled for 2-3 hours and then cooled. After addition of 100 ml water, the 9 solution was made strongly basic with a 50% solution of NaOH. The ammonia was distilled off and absorbed into a saturated solution of boric acid. The amount of borate formed was determined by titration with standard 0.1 M hydrochloric acid using the mixed indicator bromocresolgreen-methyl red. The analysis for cobalt was done by a complexometric titration with EDTA. Approximately 0.1 g complex was decomposed by heating in 30 ml concentrated sulphuric acid to dryness; the crystals being finally heated to red heat to help decompose the complex. The crystals were then dissolved in 20 ml dilute sulphuric acid and transferred to a conical flask. The pH was adjusted to 10 with concentrated ammonium hydroxide solution. Approximately 0.5 g indicator powder, Murexide mixed with sodium chloride in the ratio 1:500, was added and the solution titrated with standard 0.015 M EDTA solution. The analysis for chloride was performed by a potentiometric titration of the chloride ion with 0.01 M silver nitrate solution, using a silver-silver chloride electrode and a saturated KC1 calomel electrode as a reference. Approximately 0.2 g complex was dissolved in 25 ml of 0.25 M sodium hydroxide solution and allowed to stand for two hours. Standing for 24 hours released no more chloride ion. The solution was then brought to pH4 and 1 M ionic strength with perchloride acid and sodium perchlorate solutions and titrated to the end point of 259 mv with standard silver nitrate solution. See dE Figure 1. The end point was determined by plotting versus V, where E is the potential and V the volume of AgN03 added. 10 0 1— ____! ________ 0 2 Silver nitrate added, 4 G solution mis. 0.5 240 - 250. .260 270 . 280 Potential, mv 11 The spectrum of the product of this reaction was compared to the reported spectrum of the aquoammine and diaquo complexes (16). This comparison showed that the product of the reaction was the aquoammine complex, and therefore the release of chloride is the only reaction that takes place. Analytical Results Anal. Calculated for CoCDH^ClNHa.^HaO: Cl, 10.11; N, 19.98; Co, 16.81 . Found: Cl, 10.11; N, 19.81; Co, 16.67 . Instruments The constant temperature bath was controlled by a mercury regulator which activated heating and cooling circuits. This allowed the temperature to be controlled to within 0.02°. The 25° thermometer was calibrated against a U.S. Bureau of Standards thermometer. A Beckman Research model pH meter was used both as a potentiometer and as a pH meter. A correction curve was calibrated to allow for the effects of 1 M ionic strength on the Beckman type E (blue) glass electrodes. Ultraviolet and visible spectra were taken on a Beckman model DU quartz spectrophotometer and a Cary model 14 PM recording spectrophotometer. Infrared spectra were taken on a Perkins-Elmer model 337 spectrophotometer using potassium bromide disks. NMR spectra were taken on a Varian associates model A-60 spectrometer operating at 60 Mc/sec. The temperature of the probe was room temperature. Spectra were recorded at a 500 sec sweep rate, generally with an amplification of 25. Chemical shifts were measured 12 relative to tetramethylsilane (IMS) as external standard. The chemical shifts are expressed in terms of t (ppm). Measurement of the Acidity Constant The acidity constant of Co (0H)2C1NH3 was determined by acid-base titration using the pH meter. The pK values were obtained graphically from the linear plot of -log(HA/A~) versus pH. In this experiment a known amount of complex was dissolved in water and made up to 1M ionic strength with sodium perchlorate. A known volume of NaOH solution was added to a 50 ml aliquot of this solution and the resulting pH measured. The concentrations of HA and A~ were calculated from the pH measured, the amount of NaOH added and the amount of complex present. The pH meter had been calibrated at 1M ionic strength.* This experiment could not be carried out with high accuracy. On adding the NaOH, base hydrolysis of the complex took place. At low hydroxide concentrations the rate of reaction of. the complex was quite slow and the measured pH did not shift rapidly. The acidity constant had been found to be approximately 10~12 from a previous preparatory experiment. * Therefore measurements had to be made at pH 13 and greater. At this hydroxide ion concentration the rate of base hydrolysis became appreciable, and the meter reading changed by 0.05 pH units in five minutes. Even though readings on the meter were taken rapidly, before appreciable reaction could occur, reproducibility at the higher hydroxide ion concentrations was poor. Also the complex was not very soluble, with a maximum concentration of 0.01 M, this 13 fact limited the accuracy of the experiment. The experimental results are shown in Figure 2. The plot of -log(HA/A-) versus pH is expected to have a slope of -1. This is explained in Appendix 3. The experimental line of best fit has a slope of -1.1. -1 3 The acidity constant Kwas found to be 4.68 x 10 . The neutralization constant K equals Ka/Kw* Th® dissociation constant of water at 1M ionic strength is 1.85 x 10"11* (15). The error of the experiment was 10%. Therefore the neutralization constant equals 87.1 ± 8.7 M“*. Rate Determination The rate of base hydrolysis of Co (DH)2C1NH3 was measured by potentiometric titration of the chloride ion released. The equipment used for this titration was as follows. A Beckman Research model pH meter was used as a potentiometer. A saturated KC1 calomel electrode was used as a standard cell. A salt bridge was prepared by dissolving 3 g agar and 25 g NaN03 in 100 ml water and heating slowly until the solution was clear. The other electrode was a silver-silver chloride electrode. The end point was determined by titrating a known amount of chloride ion with standard AgN03 solution at 1M ionic strength. After each addition of AgN03 solution, no change in the potential was observed after 30 secs. The end point potential was found to be 259 mv. The AgN03 solution was dispensed with a 5 ml burette which could read to ± 0.002 ml. 14 0.8 0.6 0.4 0.2 0 - 0.4 0.6l II 12 Figure 2.--Plot of -log(HA/A ) versus pH, compared with unit slope, for acid-base titration of Co(DH)2C1NH3. [ 15 | A solution of the complex was heated for one hour at 40® at I pH 4 with 0.01 M AgNO, solution. No precipitate of AgCl formed. | ' j ? Therefore the AgN03 would not effect the titration by producing I ' | additional hydrolysis. ! For a specific kinetic run at OH- concentrations between | S _x I j 4.164 x 10 and 8.384 x 10 M, 0.185 g complex was dissolved in | 600 ml 1M NaClO^ solution. All large amounts of liquid were ! I | measured gravimetrically allowing for a small error due to buoyancy. j j The solution was thermostated at 25®. The required amount of NaOH i ■ ■ I solution was added with vigorous stirring. Samples of the solution 1 ; I were taken at various time: intervals using a 50 ml pipette with an ! ! enlarged tip to allow flow of the liquid in 7 sec. The pipette had ! jpreviously been calibrated by weighing the amount of water delivered. ! Enough perchloric acid was added to each sample to bring the solution to pH 4 which stopped the base hydrolysis reaction. A drop of Turgitol NPX detergent (Union Carbide Corporation) was added to prevent j | coagulation of the silver chloride precipitate formed in the titration. i jThe detergent also speeded equilibrium being reached at the end point. JThe sample was then titrated to the end point of 259 mv with I I 1 I standard silver nitrate solution. [ ) ' I In the runs at low hydroxide concentration, pH 11-12, the solution was buffered with 0.1 M Na3P0i,. In these cases two liters of water was made 0.1 M in NaaPOi, and 1 M ionic strength with sodium perchlorate. The pH was adjusted from the natural pH of 11.7 j j with 1 M perchloric acid to the pH of the particular run. Then J i L ! 600 ml of this solution was weighed out and the complex added, as a 16 solution, to start the reaction. At first, runs at a particular pH were made in triplicate. After getting good reproducibility runs were then made in duplicate. The plot of logCl-V/V^) versus time is linear. This is shown in Appendix 2. V is the volume of silver nitrate solution used in a particular titration. is the volume of silver nitrate solution used to titrate the maximum amount of chloride ion released in the reaction. The observed rate constant is derived from the slope of this line. Data for a particular kinetic run are shown in Table 1 and Figure 3. TABLE 1 DATA FOR EXPERIMENTAL RUN AT HYDROXIDE CONCENTRATION OF 4.164 x lO'1 M Time, sec (V), Volume of silver nitrate solution added, ml V/V. OO 1-V/V OO 110 0.165 0.151 0.849 190 1.355 0.307 0.693 280 1.781 0.404 0.596 375 2.143 0.487 0.513 435 2.421 0.549 0.451 525 2.666 0.604 0.396 630 2.984 0,676 0.324 725 3.212 0,727 0.273 870 3.493 0.791 0.209 OO 4.410 1? fO cy Q CvJ O O o Figure 3.--The plot of logfl-V/V^) versus time, for experimental values of Table 1. i TIM E, SEC. 18 NMR Spectra The NMR spectra of chloroamminebis(dimethyl- glyoximato)cobalt(III) and of the product of the base hydrolysis, aquoanuninebis(dimethylglyoximato)cobalt(III) chloride, were measured as further proof of identification and the trans configuration of the compound. For the chloroammine complex, intramolecular hydrogen bonding gave rise to a broad line at 1106 c/sec below TMS, (t = - 8.43 ppm). The width at 1/2 height was 20 c/sec. The literature reported value ' of this line was 1043 c/sec (x = - 8.47) (3). Thus the spectra measured had a frequency the same as the reported value, within experimental error. The methyl protons gave a line at 144.5 c/sec (X = 7.59 ppm) below TMS. For the aquoammine complex, lines were observed at 214 c/sec below TMS (x = 6.43 ppm) due to coordinated water and 157 c/sec below TMS (x = 7.48 ppm) due to methyl protons. This spectrum was measured in 75% DMSO and 25% acetone. The line due to coordinated water was proved to be such, by its shift due to substitution of the H2O by DMSO. The shift was slow, being 20 c/sec in 12 hours. The broad band due to the hydrogen bond of the aquoammine complex was not observable, the spectrum being measured in DMSO, DMSO + acetone, DMA and H2O. No reasons are given for this. The results for the chloroammine complex are further evidence for identification of the compound and for trans configuration. The molecular structure of the dimethylglyoxime complex is 19 shown in Figure 4. The NMR spectra are shown in Figures 5-7. Infrared Spectrum In bis(dimethylglyoximato)cobalt(III) complexes, infrared absorption bands occurring around 235G cm 1 and 1750 cm-1 are observed which are attributed to hydrogen bonding. This hydrogen bonding is said to be the proof of trans configuration of the molecule (3). The infrared spectrum was taken of the complex whose analysis showed it to be Co (DH)2C1NH3. Absorption peaks were observed at 3550 cm-1, 3410 cm-1, 3115 cm-1, 2350 cm-1, and 1750 cm-1. The first three peaks have been assigned as N-H stretching modes (3). The typical hydrogen bond stretching frequency was observed at 2350 cm-1 showing the trans character of the molecule. The peak at 1750 cm-1 was assigned as the O-H-O bending frequency. The assignment of the last two peaks was supported by frequency shifts upon deuteration t , (2, 3). The results of this spectrum is further evidence for the correctness of the elemental analysis and is said to be proof of the trans configuration of the molecule. The complete spectrum is shown in Figures 8 and 9. Absorption Spectrum of Chloroamminebis(dimethyl- glyoximato)cobalt(III) The absorption spectrum of this compound was measured by dissolving the required amount of complex to give concentration ranges between 2 x 10“ 5 M and 2 x 10“2 M. The solutions were made 1 M 20 fO X u I ro o u o K) X O fO X Figure 4.--The molecular structure of chloroammine- bis(dimethylglyoximato)cobalt(III). 21 1100 C Y C L E S /S E C Figure 5.--NMR spectrum, of the hydrogen bond of chloroamminebis(dimethylglyoximato)cobalt(III). 22 100 200 150 C Y C L E S /S E C . Figure 6.--NMR spectrum of the methyl protons of chloroamminebis(dimethylglyoximato)cobalt(III). 23 Spinning side band 200 CYCLES/SEC. Figure 7.--NMR spectrum of aquoamminebis(dimethyl glyoximato) cobalt (III) chloride. A B S O R B A N C E 0. 0 JO .20 .30 .40 1500 3000 2000 2500 4000 3500 FREQUENCY (CfVT5) Figure 8.--Infrared spectrum of chloroamminebis(dimethylglyoximato)cobalt(III). N) - P * ABSORBANCE 0.0 .20 .30 .40 .50 .60 .70 1300 1200 1100 1000 900 - 800- - 700 600 . 500 • 400 FR E Q U E N C Y (Cf^T1) ‘ Figure 9.--Infrared spectrum of chloroamminebis(dimethylglyoximato)cobalt(III). K) Cn 26 ionic strength with sodium perchlorate. All the volumetric flasks were wrapped with aluminum foil to stop any possible photolytic reaction. The spectrum was taken within five minutes of dissolving the complex. The absorption maxima were found at 217 my with a molar absorptivity of 2.84 x 101 * M_1 cm-1 and at 246 my with a I molar absorptivity of 2.48 x lo4 * M"1 cm"1. The spectra had inflec tions at 293 my, 346 my, and 465 my with molar absorptivities of 5.18 x io3, 1.46 x 103, and 7.13 x 101 M"1cm"1 respectively. The results are shown in Table 2 and the corresponding graph of molar absorptivity versus wavelength is shown in Figure 10. Absorption Spectrum of the Conjugate Base of Chloroamminebis(dimethylglyoximat)cobalt(III) Due to the fact that pK& of the compound had been found to be 12.33, the spectrum was measured at pH 14 to insure complete conversion to the conjugate base. The spectrum was measured on the Cary 14 recording spectrophotometer. The solutions were prepared by pipetting the required amounts of complex solution into volumetric flasks to which the required amounts of NaOH had already been added. The mixture was shaken up and the spectrum taken as rapidly as possible. On calculation, less than 5% reaction of the complex by base hydrolysis could have occurred. An absorption maximum was found to be at 255 my with molar absorptivity of 3.80 x io* M-1 cm-1. The results are shown in Table 3 and the corresponding graph of molar absorptivity versus wavelength is shown in Figure 11. 27 TABLE 2 THE MOLAR ABSORTIVITIES OF CHLOROAMMINEBIS(DIMETHYL- GLYOXIMATO)COBALT(III) AT VARIOUS WAVELENGTHS Wavelength, my Molar absorptivity, M"1 cm -x 200 2.10 X 10* 210 2.45 X 10“ 215 2.78 X 10“ 217 2.84 X 10“ 220 2.76 X 10“ 233 2.12 X 10“ 240 2.26 X 10“ 246 2.48 X 10“ 260 1.49 X 10“ 280 6.01 X 103 293 5.19 X 103 300 4.39 X 103 320 2.51 X 10 3 340 1.56 X 10 3 346 1.46 X 103 351 1.36 X 103 360 1.18 X 103 380 6.65 X 102 400 2.90 X 102 420 1.33 X 102 28 TABLE 2— Continued Wavelength, my Molar absorptivity, M-1 cm"1 440 7.91 X 101 460 7.14 X 101 465 7.13 X 101 480 6.73 X 10* 500 5.17 X 101 520 3.27 X 101 540 2.27 X 101 560 1.79 X 101 580 1.36 X 10l 600 9.60 29 s o 200 500 600 WAVELENGTH, mp Figure 101--The molar absorptivities of ch1oroamminebis(dimethyl- glyoximato)cobaltJill) at various wavelengths. 30 TABLE 3 THE MOLAR ABSORPTIVITIES OF THE CONJUGATE BASE OF CHLOROAMMINEBIS(DIMETHYLGLYOXIMATO)COBALT(III) AND HYDROXOAMMINEBIS(DIMETHYLGLYOXIMATO)COBALT(III) AT VARIOUS WAVELENGTHS Molar absorptivity, M-1 cm-1 Wavelength, mp Co(DH)(D)C1NH3 Co (DH)2OHNH3 220 1.23 X 10* 225 2.17 X 10* 229 1.01 X 10* 233 1.48 X 10* 235 1.09 X 10* 240 2.28 X 10* 1.34 X 10* 245 1.53 X 10* 255 3.80 X 10* 260 3.37 X 10* 1.28 X 10* 280 1.80 X 10* 8.09 X 103 300 00 • <o O' X 103 5.46 X 103 320 4.74 X 103 3.58 X 103 340 3.32 X 103 2.28 X 103 360 2.46 X 103 1.48 X 103 380 1.72 X 103 9.46 X 102 400 1.32 X 103 6.09 X 102 420 1.02 X 103 4.21 X 102 440 6.96 X 102 2.61 X 102 31 TABLE 3--Continued Molar absorptivity, Wavelength, my Co(DH)(D)C1NH3 Co (DH)20HNH3 460 4.01 X 102 1.79 X 10 480 2.13 X id2 1.22 X 10 500 1.56 X 102 8.77 X 10 520 1.06 X 102 5.99 X 10 540 7.05 X 101 4.13 X 10 560 4.36 X 101 2.82 X 10 580 2.99 X 10l 2.14 X 10 600 2.41 X 101 1.58 X 10 500 200 300 400 600 WAVELENGTH, mp Figure 11.--The molar.absorptivities. of the conjugate base.of 'chloroamminebis(dimethylglyoximato)cobalt(III) at various wavelengths. 33 Absorption Spectrum of Aquoamminebis(dimethyl glyoximato) cobalt (III) chloride The spectrum of this compound was measured by dissolving the required amount of chloroammine complex in 0.1 M sodium hydroxide solution, and allowing the base hydrolysis reaction to occur for ten half-lives. During this time the flasks were wrapped with aluminum foil to prevent any photolytic reaction. The solutions were then brought to pH 2 at 1 M ionic strength with perchloric acid and sodium perchlorate. The spectrum was measured on the Cary 14 record ing spectrophotometer. An absorption maximum was found at 245 my with molar absorptivity of 1.79 x 10** M"1 cm"1. The results are shown in Table 4 and the corresponding graph of molar absorptivity versus wavelength. This is shown in Figure 12. Comparison of these values with reported values show the product of base hydrolysis to be hydroxoamminebis(dimethylglyoximato)cobalt(III) (16). Absorption Spectrum of the Conjugate Base of Aquo amminebis (dimethylglyoximato)cobalt(III) chloride When this compound is dissolved in NaOH solution, a proton is lost from the aquo ligand forming the hydroxo complex. The pK of this ionization is 6.98 (16). Therefore if the spectrum is measured at pH 12 the compound will be 100% in the hydroxo form. The solution was prepared by dissolving the required amounts of the chloroammine complex as described in the previous section. The solutions were brought to pH 12 and 1 M ionic strength with perchloric acid and sodium perchlorate solutions. The spectra were measured on the Cary spectrophotometer. An absorption maximum was found at 245 my with a 34 TABLE 4 THE MOLAR ABSORPTIVITIES OF AQUOAMMINEBIS(DIMETHYL GLYOXIMATO) COBALT(II I) CHLORIDE AT VARIOUS WAVELENGTHS I Wavelength, mp Molar absorptivity, M-1 cm“l 220 1.12 X 0" 229 1.08 X 0* 235 1.38 X 0“ 240 1.69 X o- 245 1.79 X ; 0“ 250 1.71 X °h 260 1.42 X 0“ 280 8.98 X O3 ! 300 5.07 X 320 2.96 X o3 340 1.83 X O3 360 1.22 X 03 ; 380 8.88 X 02 400 6.62 X 02 420 4.64 X 02 440 3.35 X o2 . I 460 2.37 X o2 480 1.55 X 02 i 500 9.44 X o1 ! 1 520 5.42 X O1 1 540 3.63 X o1 560 2.26 X ol 580 1.55 X o1 600 1.11 X o1 200 300 500 400 600 WAVELENGTH, mjj Figure 12.--The molar absorptivities of aquaamminebis(dimethyl glyoximato) cobalt(III) chloride at various wavelengths. 36 molar absorptivity of 1.53 x lCf*M-1 cm-1. The results are shown in Table 3 and the corresponding graph of molar absorptivity versus wavelength is shown in Figure 13. Acid Hydrolysis The rate of acid hydrolysis was measured identically to the measurement of the rate of base hydrolysis except that no extra sodium hydroxide solution was added. The reaction was followed at 25°. The amount of chloride ion released was titrated at weekly intervals. Since the amount of complex reacting had been accurately weighed, the amount of chloride released at t = 00 was known. The volume of silver nitrate solution needed to titrate this amount of chloride ion was calculated. This value of V was used in the plot of log i ' (1-V/V^) to obtain the rate constant for acid hydrolysis. By spectral investigation it was seen to be impossible to follow the reaction at higher temperatures. At 40° the complex decomposed. The results are shown in Table 5. The calculated result for the acid hydrolysis is approximately 10-7 sec-1. * I £ ,M ~ I C M “ 37' I 400 500 300 W A V ELE N G TH , my Figure 13.--The molar absorptivities of hydroxoammine- bis(dimethylglyoximato)cobalt(III) at various wavelengths. 38 TABLE 5 DATA FOR THE ACID HYDROLYSIS OF Co CDH)2C1NH3 Time, hours (V), Volume of silver nitrate solution added, ml v/v. i-v/v„ 509 0.519 0.125 0.875 835 1.121 0.270 0.730 1181 1.821 0.440 0.560 j 4.153 (calculated) i 8 i RESULTS As previously described, the complex chloroamminebis(dimethyl glyoximato) cobalt (III) dissociates into an acid base pair. This is shown in equation (1). K Co(0H)2C1NH3 + OH" t Co(DH)(D)C1NH3 + H20 (1) The value of the neutralization constant K was found to be 87.1 ± 8.7. The rate of base hydrolysis of chloroamminebis(dimethyl- glyoximato)cobalt(III) was measured at 25° and 1 M ionic strength. This reaction is shown in equation (2) kl Co (DH)2C1NH3 + OH" ? Co(0H)20HNH3 + Cl" (2) k‘- i The reverse reaction must be considered negligible since the analysis of chloride released from the complex was within 0.5% of the expected value. The results of base hydrolysis reactions are shown in Table 6. The graph of the observed pseudo first order rate constant, versus hydroxide ion concentration is given in Figure 14. The line drawn is the computer calculated line of best fit. This line is not a straight line, within experimental error. 39 i DATA FOR THE BASE TABLE 6 HYDROLYSIS OF Co (DH)2C1NH3 pH read from pH meter Corrected pH Hydroxide concentration, M Buffer k Run 1 , , sec-1 obs Run 2 Run 3 Average k , sec"1 obs, — — 4.164*10-1 none 1.85*10"3 1.82x10"3 1.87X10-3 1.85X10"3 — — 2.096x10“1 none 9.49x10”* 9.18*10-* — 9.33*10'* — — 8.384X10-2 none 3.77X10-* 3.70*10"* 3.73*10-* 11.300 12.147 7.59 xio-3 Na3P0v+HC10i, 4.37*10-5 4.50*10-5 4.49*10-5 11.020 11.861 3.93 xio-3 Na3POi,+HC10i, 2.48*10-s 2.51*10“s — 2.50*10"5 10.650 11.457 1.55 xio-3 Na3P0^+HC10i, i.oi*io-5 1.04*10"5 1.03*10"5 -fc. © L O G O F OBSERVED RATE CONSTANT, SEC." 41 pOH Figure 14.--Graph of log of the observed rate constant versus pOH, for base hydrolysis of chloroamminebis(dimethyl glyoximato) cobalt (III). DISCUSSION From Figure 14 the base hydrolysis of Co (DH)2C1NH3 appears to occur by a mechanism which has two parallel reaction paths. One path is first order in hydroxide ion concentration and first order in complex concentration. The other path is second order in hydroxide ion concentration and first order in complex concentration. The term which gives the second order hydroxide ion dependence becomes fifst order dependent at high hydroxide ion concentrations. This together with the knowledge of the acid-base equilibrium of the complex gives an observed rate law as shown in equation (3). 4 [Cl 1 _ r a [OH ] + b[OH ]2 ] j - c ] dt 1 + c[OH“] t In equation (3) a, b, and c are constants. Ct is the total complex concentration. The experimental data were analyzed by a non-linear least squares computer program written by R. H. Moore and R. K. Zeigler at Los Alamos Scientific Laboratory, Los Alamos, New Mexico. The data were analyzed in two different ways. Firstly the computer analyzed the data such that it was free to find the best value of all parameters. This result is shown in Table 7. Then the data were analyzed with the value of c fixed. The constant c was assumed to be the neutrali- i zation constant. This is explained in detail later in the discussion. These results are shown in Table 8. 42 43 TABLE 7 NON-LINEAR LEAST SQUARES ANALYSIS OF DATA, ALL PARAMETERS FREE Hydroxide Concentration, M k , , sec“* obs Computer calculated k , , sec"* obs Percentage Deviation 4.164x10"1 1.85x10-3 1.816xl0“3 + 1.83 % 2.096X10-1 9.33x10-** 9.268x10"** + 0.652% 8.384xl0“2 3.73x10-** 3.848x10-** - 3.15 % 7.59xl0-3 4.50x10"5 4.459x10“5 + 0.874% 3.93x1b-3 2.50x10“5 2.475xl0-5 + 1.05 % 1.55x10“3 1.03X10'5 1.041X10"5 - 1.10 % TABLE 8 NON-LINEAR LEAST SQUARES ANALYSIS OF DATA, K AT THE EXPERIMENTAL VALUE OF 8.71x10* M"* FIXED Hydroxide Concentration, M k . , see"* obs* Computer calculated k0bs- sec"' Percentage Deviation 4.164x10“* 1.85X10-3 1.810X10-3 + 2.19 % 2.096x10-* 9.33x10"** 9.256x10"** + 0.807% 8.384X10"2 3.73x10-** 3.859x10-** - 3.44 % 7.59X10-33 4.50xl0"s 4.489xl0"5 + 0.198% 3.93x10"3 2.50x10“5 2.477x10-5 + 0.941% 1.55X10"3 1.03x10"5 1.034X10-5 - 0.443% i < 44 The experimental value of the neutralization constant is 87.1 M"1. The Computer'results, all parameters free, are as follows: The value of c is 1.026 x lo2 M"1 with standard deviation of 3.747 x 101 M"1. The value of a is 7.098 x 10”3 M“2 sec"1 with standard deviation of 2.347 x 10“3 M"2 sec"1. The value of b is 4.409 x 10"1 M"2 sec-1 with standard deviation of 8.400 x 10"3 M"2 sec-1. The computer results, where c was held at the experimental value of 87.1 M"1, are as follows. The value of a is 6.998 x 10-3 M"2 sec"1 with standard deviation of 1.086 x 10-1* M"2 j sec-1. The value of b is 3.723 x 10"1 M~2 sec"1 with standard deviation of 5.237 x 10-3 M“2 sec"1. Acid-base Characteristics of dimethylglyoxime Complexes In dimethylglyoxime complexes of cobalt (HI) there is an acidic proton on a dimethylglyoxime ligand. The acidity of this proton is dependent on the additional ligands on the complex. For example, if one of the ligands is the aquo ligand, a proton is only lost from this aquo ligand. The pK of this dissociation is 6.98 for the aquoammine complex (16). Excess NaOH does not remove a proton from a dimethylglyoxime ligand in the aquo complex. The proton removal is always reversible. In the case where the additional ligands are ammonia or pyridine or chloride, the proton is then removable from a dimethylglyoxime ligand. These additional ligands have protons of weaker acidity than the dimethylglyoxime ligand, or they have no protons as in chloride. 45 These facts are deduced from spectral changes (16) and pH titrations. When a proton is lost from a dimethylglyoxime ligand, there is a shift of the peak at 245mp to 255 mp. This is a reversible change. In aquo complexes, there is no spectral change on addition of excess sodium hydroxide. Thus when Co (DH)2C1NH3 is dissolved in sodium hydroxide solution, there is a dissociation of a proton from a dimethylglyoxime ligand forming a conjugate acid-base pair. This dissociation has been found to have a pK of 12.33. This dissociation has also been seen r a to be reversible. Equation (4) shows the equilibrium where C is the & - conjugate acid and the conjugate base of Co (DH)2C1NH3. represents the molecular formula Co(DH)(D)C1NH3 where D is the jdeprotonated dimethylglyoxime ligand. K is the equilibrium constant. K Ca + OH' t Cb + H20 (4) Therefore by definition K =....CS) [Ca][0H~] Also the sum of the concentrations of the conjugate acid and conjugate base equals the total complex concentration as shown in equation (6). [Ct] » [Ca] ♦ [Cb] (6) Solving equations (5) and (6) for C and C, : 3. D 46 I [C ] [CJ - -- S-z- (7) a 1+K[0H ] K[OH"][C ] [C ] = ------- -L_ (8) D 1+K[0H ] There is also the possibility of a tautomeric equilibrium as shown in equation (9) k.i Co(DH) (D)C1NH3 £ Co (DH)2C1NH2 (9) The equilibrium constant of this tautomerism is unknown but assumed to be small. The reaction is a transfer of a proton from an ammonia ligand to a dimethylglyoxime ligand. 3+ The rate of base catalysed hydrogen exchange of Co(NH3)6 ! is 1.6 x 10s M-1 sec-1. All cobalt ammine complex exchange at approximately the same rate (18). Since this is such a rapid exchange, it could be assumed that the rate of proton transfer in equation (9) is also rapid, at least more rapid than the relatively I slow base hydrolysis reaction. The equilibrium would be expected to | be far to the left since the protons on the dimethylglyoxime ligands are more acidic than the protons on the ammonia ligand. Thus in basic solution there are assumed to be three possible species present, all in equilibrium with each other. These are C , Co (DH)2C1NH3; C, , Co(DH) (D)C1NH3; and C. ., Co (DH)2C1NH2". * SL D tEUu Mechanisms of Reaction The data, with computer analysis, supports the statement that the compound chloroamminebis(dimethylglyoximato)cobalt(III) reacts in a base hydrolysis reaction by a mechanism which has two parallel pathways. One pathway is first order in hydroxide ion concentration and the second path second order in hydroxide ion concentration. Both paths are first order in total complex concentration C^.. The observed rate law is shown in equation (3). By the experimental methods used, this rate law is the rate of chloride release. Spectral analysis of the product of the reaction shows that this rate is also equal to the rate of base hydrolysis. Equation (3) may then be rewritten d[C ] a[OH~] + b[0H"]2 - — — = [ H C J do) dt 1+c[OH-] Therefore the experimental pseudo first order rate constant for the base hydrolysis of Co (DH)2C1NH3 is of the form a[0H“] + b[OH~] 2 kobs = [-----------:-----1 (11) ODS l+c[0H ] The following scheme shows the possible reaction pathways for the base hydrolysis reaction. 48 ki* (OH*) k2 (OH~)______ Co (DH) (D)NHi--- ► JUl JjL kl(OH*) kUOH*)2 taut kI The rate constants identify the pathways. In the scheme the complete reactions are not shown. The complete reactions will be shown in detail as each path is discussed. In this discussion section the mechanism of the base hydrolysis reaction will be discussed in terms of the following rate laws. These rate laws all fit the experimental rate data. The rate laws are: d[ct ] dt = k2[OH‘][Cb] ♦ ki[Cb] ( 12) d[C.] — * k2[OH-][Cb] ♦ k'[Ctaut] dt (13) d[C£ dt - k ' [0H-][c taut] ♦ kx[cb] (14) d[C^ dt (15) The reaction paths will be discussed such that the paths zero in hydroxide ion concentration, k; and kj, will be compared. Also the paths first order in hydroxide ion concentration, k2 and k£, 0 49 will be discussed and compared. From this discussion the rate law which has the most precedent will be chosen from equations (12) to (IS). All the pathways are to be discussed in terms of an S^ICB mechanism. Paths k,^ and kj, which are considered to be S^2 type mechanisms will not be discussed for reasons given in the introduction. If Ca is considered to react in an S^ICB mechanism through path k4, the amido complex formed by removal of a proton is Co (DH)2C1NH2. This is identical with the species ^ ut» which reacts through path kj. Thus an S^ICB mechanism of Ca reacting is kinetically equivalent to Ctaut reacting through path k(. Discussion and Comparison of Pathways k2 and k^ The path k2 of the previous scheme may be described by the following equations. K2(fast) Co(DH)(D)ClNH^ + OH" t Co(DH) (D)C1NH-2 + H20 (16) k2 Co(DH) (D)C1NH“2 * Co(DH)(D)NH2 + Cl" (17) (fast) Co(DH) (D)NH2 + 2H20 Co (DH)2OHNH3 + OH" (18) The path k^ may be described by the following equations (fast) Co (DH)2C1NH2 + OH" * Co(DH) (D)C1NH£2 + H.O (19) k2 Co(DH)(D)ClNH:2 - > • Co(DH)(D)NH" + Cl* (20) (fast) Co(DH)(D)NH“ + 2H,0 + Co(DH)20HNH, + OH- (21) Comparing equations (17) and (20) it is seen that the reactive intermediates are identical. The difference in these paths is only in the order of generation of the intermediate. Thus the mechanisms of paths k2 and k| are kinetically equivalent. The mechanisms of these paths are postulated as being S^ICB. This mechanism, as previously explained, has precedent in Co(III) reactions. Therefore k2 and kj are likely reaction pathways. Since they are equivalent, the reaction will be considered in terms of path k2 only. ! { The rate of reaction of this pathway is then equal to ikJCo(DH) (D)C1NH;2]. Discussion and Comparison of Pathways kt and kj The path ki is shown by the following equations ki Co(DH)(D)ClNH; Co (DH)(D)NH3 + Cl" (22) (fast) Co (DH)(D)NH3 + H,0 - * ■ Co (DH)20HNH, (23) Equation (22) is the rate determining step with rate constant | |kj. This step is independent of hydroxide ion concentration. It has jbeen experimentally determined that the value of k3, the acid 51 I i ' hydrolysis rate constant is approximately 10~7 sec"1. The value of kx, if calculated from the experimental value of constant a of equation (11) is 6.917 x 10-5 sec-1. Therefore for the mechanism of equations (22) to (23) to be correct, removal of a proton from C £ L to form must increase the rate of hydrolysis, independent of hydroxide ion, by three orders of magnitude. This conclusion, in the absence of further data, would make this pathway an unlikely reaction route. When is formed from C& a proton is lost from a dimethylglyoxime ligand. This leaves an electron pair on the oxygen atom. Refer to Figure 4. In cobalt amido complexes the electron pair on the nitrogen is thought to form an intermediate, stabilized by tt bonding to the adjacent cobalt atom. The formation of this intermediate aids in the loss of the halide ion (19). In the dimethylglyoxime complex the oxygen atom is not adjacent to the cobalt i ' atom. A transfer of electrons through the nitrogen to the cobalt i I cannot automatically be assumed to take place. Therefore in the absence of further information this is probably unlikely. Therefore i one cannot assume that path ki reacts through an S^ICB mechanism. The path kx is concluded to be an unlikely reaction route. t The path k’ may be described by the following equations. ! k * I Co (DH)2C1NH; Co (DH)2NH2 + Cl" (24) I j Co (DH)2NH2 + H20 - » ■ Co (DH)2OHNH3 (25) ! I As described previously the species C^, and C u are all assumed to be in equilibrium with each other in basic solution. Thus the reaction to form C by loss of a proton is shown in t&ut equation (26). (f3-St) Co(DH)2C1NH3 + OH' t Co(DH)2ClNHi + H20 (26) Combining equations (26), (24), and (25) gives the series of equations necessary for an S^ICB mechanism. Since it is kinetically impossible to say by which series of equilibria ctaut was formed, equations (24) and (25), of which (24) is rate determining, are equivalent to an SN1CB mechanism. Thus the path kj is a reasonable reaction route. It is essentially zero order in hydroxide ion concentration. The rate is then equal to kj[C a 3 or equivalently k'[Co(DH)2ClNHi]. Comparison of the Observed Rate Constants to Calculated Rate Constants Referring to Figure 14, one could possibly draw a straight line through all the data points. However, considering the experi mental error of the points, the line of best fit is a curve. A straight line assumes a mechanism with only one path, a path first order in hydroxide ion concentration and first order in complex concentration. It has been established from the acid-base titration that the compound exists as a conjugate acid-base pair. t | If only C^ reacts with hydroxide ion, then the observed rate i |would drop rapidly as the pH of the reaction was changed from 12 to i J 11. The rate in this case_would be first order in hydroxide JLon___ 53 concentration and first order in concentration. This rate law is represented by equation (27). d [ c n dt = k2[0H][Cb] (27) Substituting equation (8) into equation (27) and integrating gives a rate constant as shown in equation (28). "1 2 k,K[OH ] k = ------------ (28) % 1+K[0H~] If the reaction were just the reaction of C , then by a a similar calculation, substitution of equation (7) into equation (29) gives equation (30). d[Cl ] ------ = k [0H'J[C ] (29) dt a k [OH"] k = : - (30) a 1+K[0H'] The observed rate constants and the calculated contributions to this rate constant of and are shown in Figure 15. Considering the form of equations (28) and (30) it would be unlikely that the sum of these two rates were a straight line. Thus it is false to assume that the data lead to a mechanism with only one reaction pathway. L O G O F R A TE CONSTANT, SEC.- 54 (V pOH Figure 15.--The log of the observed constants and the calcu lated contributions of Ca and to'this rate, versus pOH. _ _ _ _ observed rate constant contribution to the rate of CT -------------------------- -a . . contribution to the rate of 55 As discussed earlier the path k * Cctaut-3 *s preferred to path kjtC^]. Also paths k2 and k£ are equivalent. The equilibria for the formation of and Co(DH) (D)ClNHi are evidently far to the left. Therefore t W - K > c cbi < 3l> [Co(DH) (D)ClNHj2] * K2[0H][C.] (32) The rate equation which fits the data and has the most precedent is shown by equation (33). d[Ct] = [Ctaut] + k2[Co(DH)(D)ClNH“2] (33) dt Substituting equation (32) into equation (33) gives a rate law of the same algebraic form as the proposed rate laws previously given. Then substituting equations (31) and (32) into equation (33) gives equation (34). d[Cj _ = k. Kl[Cb] + k2 K2[OH"][Cb] (34) dt Substituting equation (8) for and integrating gives an observed rate constant as shown in equation (35). k' K K[OH“] + k K K[0H"]2 k . = ---- -------------?— -------- (35) 0bS 1+K[0H-] 56 This equation is of the same algebraic form as equation (11), the experimentally observed rate constant. From equation (35) it is seen that the constant < ? equals the neutralization constant K. Therefore the earlier assumption that c = K in the computer analysis was valid. ( The computer analysis results for the constants kj Kx, k K and K are shown in Tables 9 and 10. 2 ^ TABLE 9 COMPUTER RESULTS FOR RATE AND EQUILIBRIUM CONSTANTS, ALL PARAMETERS FREE Constant Value of Constant Standard Deviation Experimental | Value K 1.026 x io2 M-1 3.747 x 101 M'1 8.71 x 101 M“l ! i k; k, 6.918 x IO"5 M"1 sec-1 2.288 x 10~s M_1 sec-1 — O l 4.297 x 10"3 M"1 sec'1 8.187 x IO"5 M"1 sec-1 — < J 2+ The rate constant for acid hydrolysis of Co (NH3)sC1 is < 1.71 x 10“6 sec-1 (20). The rate constant for base hydrolysis is | 0.85 M-1 sec"1. j The hydrolysis rates for chloroamminebis(dimethyl- | glyoximato)cobalt(III) are at least two orders of magnitude slower than corresponding results for the pentaammine complex. TABLE 10 COMPUTER RESULTS FOR RATE AND EQUILIBRIUM CONSTANTS, K HELD AT EXPERIMENTAL VALUE OF 8.71 x IQ1 M*1 Value of Standard Constant Constant Deviation 8.035 x io-5 M“l sec-1 1.247 x io-6 M“l sec"1 4.274 x 10"3 M_I sec"1 6.013 x IQ-5 M"1 sec"1 The acid hydrolysis rate constant for Co (DH)2C1NH3 is approximately 10“7 sec-1. The base hydrolysis results are shown in Tables 9 and 10. The possible reasons for the slowness of the hydrolysis of the complex Co(DH)2C1NH3 could be due to differences in the values of j Kj and K2 compared to the corresponding equilibrium constant of i the pentaammine complex forming the amido intermediate. i i In the dimethylglyoxime complex the intermediate Co(DH)(D)C1NH~2 has charge -2. It is formed by loss of a proton | from a molecule of charge -1. In the pentaammine complex the molecule j f jis charged +2. The reactive intermediate of the SN1CB mechanism jis charged +1. Therefore by a purely electrostatic argument, r | removal of a proton to form the pentaammine intermediate would be 1 easier. The base hydrolysis of the pentaammine complex would then be expected to be more rapid. In the absence of further data no reasons kj Kx 58 can be given for this difference in rates. Further work, it is felt, should be done on the base hydrolysis of complexes of the form Co (DH)2C1Y where Y is a ligand with no acidic protons to see how this effects the rate of base hydrolysis. I I I APPENDIX I As previously stated, the original aim of this project was to prepare the aquoamminebis(dimethylglyoximato)cobalt CHI) cation to continue work being done in this laboratory on the anation of various cobalt dimethylglyoxime complexes. The first method of preparation was the acid hydrolysis of chloroamminebis(dimethylglyoximato)cobalt(III). This reaction was found to be much too slow. After heating a solution of the complex, approximately 0.01 M, at 40° for five hours, no precipitate of silver chloride was formed on adding silver nitrate solution. If the same solution was heated at a higher temperature, the compound decomposed. This was determined by spectral investigation. The second method of preparation involved heating a 0.01 M solution of the complex, with equimolar quantities of silver nitrate. After heating for two hours at 60°, no precipitate of silver chloride was seen. If the temperature was raised, the compound decomposed. To explain what is meant by these compounds decomposing, it ! t must be stated that all dimethylglyoxime complexes of cobalt(III) j appear to have a peak in their ultraviolet spectra at approximately 245 mp. When the spectra of the compounds were measured after ■ heating the peak at 245 mu, assumed to be due to the dimethylglyoxime i j ligands, broadened and eventually disappeared. The products were l unknown,' and no chloride ion was released. 1 59 The next preparative method involved heating the chloroammine complex with silver oxide. 2.3 g of the chloroammine complex were heated with 1.4 g silver oxide in 50 ml water at 50° for four hours. The mixture was filtered to remove the silver chloride formed and unreacted material. The dark brown solution was evaporated under reduced pressure at a maximum temperature of 60° until almost dry. On I addition of ethanol to this solution, a small amount of pale brown, flocculent crystals precipitated out. If the water solution was evaporated to dryness, a second product, dark brown shiny crystals, was formed. Both these products had similar ultraviolet spectra with peaks at 243 my. The dark brown crystals, the major product, were recrystalized by dissolving in ethanol and precipitating with ether. After repeating this procedure three times the spectrum had a peak at 243 my with an extinction coefficient of 1.71 *10** M”1 cm-1. This I extinction coefficient is based on the molecular weight of the hydroxoammine complex. Weight loss on heating showed the complex to have three waters of hydration. | The compound was sent for analysis three times. Each sample :was prepared the same way except that the temperature of evaporation was changed from 60° to 50° to 40° in the three cases to stop any decomposition due to heating. The first two analyses were sent to Elek Microanalytical Laboratories, Torrance, California 90502. The third analysis was sent 1 I to Alfred Bernhardt Mikroanalytisches Laboratorium, 433 Muehlheim I | (Ruhr) Germany. 61 I l f Analytical Results Anal. Calculated for CoCDH^HzOMVSH 0: C, 25.4; N, 18.6; H, 6.1. Found: (1) C, 26.6; N, 18.3; H, 5.7. (2) C, 26.5; N, 17.9; H, 5.4. (3) C, 25.03; N, 17.27; H, 5.35. In each case the analyzed results were not satisfactory in that the carbon to nitrogen ratios were always lower than the calculated value. The aquoammine complex was allowed to react with various ions to see if the spectral change on reaction was great enough to follow the reaction on a spectrophotometer. These ions were: SCN”, N^”, Cl”, Br”, I”, and CN”, and also pyridine. The concentrations of all these materials were 0.1 M. Also 6 M SCN” ion was tried. In all i • i | these cases there was no change in the spectra with up to 24 hours i j reaction at 40°. In these experiments the spectra were only measured j in the ultraviolet region. Then on the possibility that the compound was a polymer that i formed either during preparation or recrystalization, molecular weight | studies were carried out. > | The molecular weight measurements were carried out on a j Mechrolab Model 301A Vapor Pressure Osmometer. This instrument works on the principle that a solution will always have a lower vapor pres- I sure than the pure solvent. A drop of solvent and solution are ! j suspended side by side in a thermostated chamber. Because of the •difference in vapor pressure, of the drops, a differential mass 62 transfer will occur between the drops causing a temperature difference. This temperature difference is proportional to the vapor pressure lowering and therefore to the solute concentration. The best results are obtainable from organic solvents due to the greater rate of evaporation. Examples of these solvents are benzene and alcohol. To use these solvents the thermostated chamber must be scrupulously cleaned of other solvents. Thermistor probes are available for water and organic solvents at 37° and 25°. All the readings were done using the aqueous probe since this was the only probe available at the time. The machine was used exactly as described in the manual. The only problem with its use was in making sure that the water used as solvent was from a constant source. Errors were found when the solvent source was changed. A calibration curve was made using NaCl solutions from | 0.01 M to 0.1 M. At lower concentrations reproducibility is greatly reduced. At much higher concentrations, condensation will t cause concentration shifts. | The machine gives results that are fairly reproducible if the f 1 , operator has his own personal method of operation of the machine l i standardized. The machine works using a Wheatstone Bridge to measure i changes in temperature of the drops. Results are given in Table 11 and Table 12. i i Typical results are given for sodium chloride solution in i jwater: 63 TABLE 11 CALIBRATION OF THE VAPOR PRESSURE OSMOMETER WITH NaCl Concentration of sodium chloride, M Experimental results, ohms Run 1 Run 2 Run 3 0.01 1.56 1.52 1.54 0.02 3.09 3.01 3.09 0.04 5.72 5.70 5.68 0.06 10.06 10.06 10.04 0.08 11.83 11.79 11.71 0.10 14.69 14.73 14.75 t j A similar calibration curve was also drawn up for sucrose, a non I ionic molecule. i | The vapor pressure lowering was then measured for the i I compound thought to be hydroxoamminebis(dimethylglyoximato)cobalt(III). i j The concentration of the complex was 0.1 M assuming the molecular ! weight for the hydroxoammine complex, which is 323. The Vapor I ; Pressure Osmometer gave a molecular weight of 1380 +_ 70, assuming a i 1 neutral molecule. Thus at some time during the preparation or the j recrystalization the compound polymerized, the polymer having a very i similar elemental analysis to the hydroxoammine complex. I ' The method of molecular weight measurement by the Vapor 1 I ; Pressure Osmometer is satisfactory, however one cannot expect accuracy 64 1 i I i 1 of better than 10% unless the operator has practiced with the machine I to improve his experimental technique. The highest accuracy expected j is 5%. The crystals thought to be hydroxoamminebis(dimethyl- glyoximato)cobalt(III) were then run through a cationic ion exchange column (Dowex 50W-X4) in the Na+ form. Four bands were observed on elution with 0.1 M NaOH solution. The hydroxoammine complex is a neutral molecule, and therefore would not be expected to stick to the column. Thus the product must have been a mixture of at least four compounds. This mixture could not be separated since the components of the mixture had very similar solubility in all the < solvents tried. The solvents were water, ethanol, ether, and j i 1,4-dioxane. The molecular weight result is thus meaningless since the compound is a mixture consisting of ions of unknown charge. ' TABLE 12 CALIBRATION OF THE VAPOR PRESSURE OSMOMETER WITH SUCROSE Concentration Experimental results., ohms of sucrose, M Run 1 Run 2 Run 3 0.04 3.15 3.12 3.10 0.06 4.65 4.64 4.59 0.08 6.14 6.10 6.09 0.10 7.65 7.63 7.58 65 Then it was observed that the rate of base hydrolysis of chloroamminebis(dimethylglyoximato)cobalt(III) was rapid enough to use this reaction as .a method of preparation of the hydroxoammine complex. Therefore 1 g of the chloroammine complex was dissolved in 50 ml 0.1 M NaOH and allowed to react for five hours. The solu tion was adjusted to pH 4 with perchloric acid and sufficient silver nitrate solution added to precipitate the released chloride ion. The solution was filtered and then passed through a cationic ion exchange column (Dowex 501V-X4) in the Na+ form. The product was then the aquoammine complex with charge +1 and should stick to the resin. This it did. The column was eluted with 0.05 M NaOH solution. The product was converted to the hydroxoammine complex, a neutral I molecule, and it passed through the column. The eluent was neutralized with perchloric acid and evaporated under vacuum, at 50°, to dryness. This preparation was carried out twice under identical conditions. The first product was insoluble in alcohol, and the | crystals were yellow. The second time the preparation was carried j out the product was found to be soluble in alcohol and these crystals were brown. The spectra of these two products were similar in that I ! they both had maxima at 245 my. ! I It was decided to study the rate of base hydrolysis of i ichloroamminebis(dimethylglyoximato)cobalt(III). None of the I preparations was successful possibly due to the fact that the compound ; polymerized under the preparative conditions used. I APPENDIX II Assume that the compound B reacts to form products C and D. Then the equation of reaction is: k B -► C + D (36) Therefore: § * * B (37) or: - kdt (38) At any time in the reaction the total complex concentration is constant. Therefore: B + C « B (39) o where Bq is the initial concentration of B. Substituting equation (39) into equation (38) gives equation (40) dC = kdt (40) B -C o Integrating equation (40) as shown in equation (41) gives equation (42) C=C t=t / = / kdt (41) C=0 o t=0 66 67 - log (1 - C/Bq) = kt (42) Bq equals the initial concentration of the complex before reaction, and also equals the total amount of chloride released. The i amount of chloride ion equals V the volume of silver nitrate solution used in the titration. Therefore B = V and C = V. o 0 0 Thus equation (42) may be written: " - log (1 - V/VJ = kt (43) . From equation (43) a plot of log (1 - V/V^) versus t is linear for a pseudo first order reaction. The slope of the line gives the observed rate constant k. APPENDIX III Assume that the compound HA dissociates to H+ and A as shown in equation (44) HA £ H+ + A" (44) f I Therefore the equilibrium constant is: j [H 1[A ) K --------- (45) a [HA] I Rearranging equation (45) [HA] log K » - log - pH (46) 3 [A"] Therefore in a plot of -log(HA/A~) versus pH, when -log(HA/A~) = 0, logK = -pH. This plot is shown in Figure 2. From equation (49), it is seen that this plot should have a slope of -1. REFERENCES 1. R. Tsuchida, M. Kobayashi, A. Nakamura, Chem. Soc. Japan, 11 38 (1936). j j 2. R. Bline, D. Hadzi, J. Chem. Soc., 4536 (1958). I ; 3. R. D. Gillard, G. Wilkinson, J. Chem. Soc., 6041 (1963). 4. A. V. Ablov, N. M. Samus, M. S. Popov, Doklady Akad. Nauk. SSSR (Russ) 106, 665 (1956). 5. A. V. Ablov, N. M. Samus, Doklady Akad. Nauk. SSSR (Russ) 113, 1265 (1957). i 6. L. A. Tschugaeff, Ber., 39, 2692 (1906). 7. R. Tsuchida, Chem. Soc. Japan 12, 83 (1937). 8. A. V. Ablov, M. P. Filippov, Russ. Jour. Inorg. Chem. 1^, 119 (1958). 9. A. Nakahara, Chem. Soc. Japan 27, 560 (1954). 10. A. V. Ablov, M. P. Filippov, Russ. Jour. Inorg. Chem. 4^ 1004 (1959). 11. A. W. Adamson, F. Basolo, Acta Chem. Scand. £, 1261 (1955). 12. R. G. Pearson, H. H. Schmidtke, F. Basolo, J. Am. Chem. Soc. 82, 4434 (1960). 13. M. Green, H. Taube, Inorg. Chem. 2, 948 (1963). 14. G. E. Dolbear, H. Taube, Inorg. Chem. 6, 63 (1967). 15. B. I. Na!>ivanets, Zhur. Neorg. Klim. 7, 417 (1962). 16. A. V. Ablov, B. A. Bovykin, N. M. Samus, Russ. Joum. Inorg. Chem. 12, 978 (1966). 69 70 17. A. V. Ablov, N. M. Samus, A. Bologa, Russ. Jour. Inorg. Chem. 8, 440 (1963). 18. F. Basolo, R. G. Pearson. Mechanisms of inorganic Chemistry (Second ed.; New York, John Wiley and Sons, Inc.), pp. 185. 19. F. Basolo, R. G. Pearson, ibid., pp. 186. | 20. F. Basolo, R. G. Pearson, ibid., pp. 164. I | 21. D. A. Buckingham, I. I. Olsen, A, M, Sargeson, J. Am. Chem. SOc. 88, 5543 (1966) 22. D. A. Buckingham, I. I. Olsen, A. M. Sargeson, J. Am. Chem. Soc. 89, 5129 (1967).

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Creator Fredericks, Michael (author)

Core Title A kinetic study of the base hydrolysis of chloroamminebis (dimethylglyoximato) cobalt (III)

Degree Master of Science

Degree Program Chemistry

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

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Advisor Wilmarth, W.K. (committee chair), Adamson A.W. (committee member), Warf, James C. (committee member)

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c17-795340

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