Thesis-1967-H315t.Pdf (3.606Mb)

A THERMODYNAlfiO STUDY OF THE AQUOOOMPLEX ,· FORMATION OF OOBALT(II) AND NICK.EL(!!) PERCHLORATE IN 1-BUTANOL By Phillip Carroll Harris " Bachelor· of Scien~ce . ,Oklahoma State University . Stillwater., Okl~oma · 1965 Submitted to the faculty of the Graduate College of the Oklahoma· State University in partial fulfillment of the requirements fo:r ·the degree of MA.STER OF SCIJSNCE · lVIay., 1967 _,_ ~Htl~ Sl"AT£ -UNM~W.!~ L. t tBR,).\~ ~ ·\INl lQ 1968 ·i, A TH.ER.WDYNAlMllC STUDY OF TH.E Ai UOCOMPLEX FORMATION OF COBALT{II) AND NICK..EL(Il) ' . P.EaCHLORATE IN 1-BUTANOL Thesis 4,pproved:,, B58B15 ii ACKNOWLEDG:MENT I wish to express my apprec:i,.ation to Dr. Tom E. lVbore for his guidance throughout the research and a!:3sistance with reading the :manuscript. I also wish to thank Dr. L. P. Varga for hi.s counsel on computational problems. I wish to thank my parents for their encouragement. This study was supported in part by funds made available by the Atomic Energy Commission, Contract No. AT(ll-1)-1175. iii TABLE OF CONTENTS Chapter Page I. INTRODUCTION • • • • • • • • • . • • 1 Background to the Problem •• • • • . " l II. THEORY OF THE PROPOSED METHOD. • • • • • . ., . 6 Corresponding Solutions ••••• • • • 6 Formation Constants and Enthalpies. 10 Assumptions of the Theory .. • • • • • • • • 13 III. APPARATUS AND PROCEDURES • • • • • • • • • 0 17 The Calorimeter • • • • • • • • • • • • • • 17 Drying Apparatus. • • • • • • • • • • • • • 2 5 Drying Pro eedur e. • • • • • • • • • • • 2 6 Thermometric Titration Procedure. • • • • • 28 IV. EXPERllVJENTAL AND TREATMENT OF DATA • • • • . 33 Solutions Used.. • • • • • • • • • • • • • • 33 Calorimetry Data. • • • • • • • • • • • • • 33 Sample Calculation. • • • • • • • • • • • • 34 Constants of the Systents. • • • • • • • • • 47 Conductivity Studies. • • • • • • • 63 N MR Studies • • • • • • • • • • • • • • • • 65 Color of the Solutions. • • . • • • • • • • • 68 V. DISCUSSlON OF RESULTS AND CONCLUSIONS. • • 0 69 Aquo complex .B'orma tion .. • • • • • • • • • • 69 lVlechanism of Aquocornplex F'ormation. • • • • 73 Reliab;i..lity of the Therm:,dynamic Quantities 79 VI. SUMMARY AND SUGGESTIONS FOR FURTHER STUDY. 82 BIBLIOGRAPHY. • • • . ' • • 0 Ill' . 85 GLOSSARY ••••• . 87 iv LI ST OF TABLES Table Page I. Heat of .Mixing of Water and 1-Butanol. • • • • • 36 II. Experimental Data and Calculated Heats of Reaction for Thermometric Ti trations. • • • • • • • • • 39 III. Least Squares Curve Fit for ii and (L]. • • • • 50 IV. Pseudo Formation Constants •• • • • • • . 57 v. Calculated E'or:mation Constants • • • • • • • • • 58 VI. Exothernrl.c Overall Enthalpies of Aquocomplex Formation (Calories per lVble) •••••..•• • • 62 VII. Revised Formation Constants, Overall Enthalpies and Entropies of Aquocomplex Formation. • • • 64 V LIST OF FIGURES Page 1. The Corresponding Solutions Method of Determining ii and [L] • • • • • • • • • • . • • • • • • • • • ll 2. Typical Formation Function of Bjerrum ~Vhere ii Approaches· N • • • • • • • • • • • • • • • • • • ll 3. Internal Parts of the Calorimeter ••••• • • • • 20 4. A Schematic of the Temperature Sensing and Heat Capacity Calibration Oircui t:cy. • • • • • • • • • 23 5.. Solution Drying Apparatus. • • • • • • • • • • • • 27 6. Densities of Cobalt(II) and Nickel(II) Perc};).lorate Solutions in 1-:Su tanol • • • • • • • • • • • • • 29 7. Typical En thalpograms. • • • • • • • • • • • • • • 32 8. Heat of 1-'Jixing of Water with 1-Butanol • • • • • • 35 9. Heat of Reaction of Water with Co ( 0104)2 • • • • • 37 10. Heat of Reaction of Water with Ni ( 010 4 ) 2 • • • • • 38 11. Corresponding Solutions Treatment fo:r 00(0104 )2• • 48 12. Corresponding Solutions Treatment for Ni(0104 )2• • 49 13. Formation Function Curve for Co(c104 )2 • • • • • • 51 14. Forrna tion Function Curve for Ni ( 0104 ) 2 • • • • • • 52 15. Pseudo lformation Constant Curve for 00(0104 )2 Aquo co mpl ex,e s. • • • • • • • • • • • • • • • • • 55 16. Pseudo Formation Constant Curve for N;i.(0104 ) 2 Aquo complexes. • • • • • • • • • • • • • • . 56 17. Heat of. Reaction of Co ( 0104 )2 with H20 as a Function of Average Ligand Nuxnber • • • • • • • • • • • • 59 18. Heat of' Reaction of Ni(Ol04) 2 with H20 as a li1unction of Average Ligand Number • • • . • • · • • • • • t • 60 Vi LIST OF FIGURES (Continued) .F'igure · Page 19. Conductivity of Co ( 0104 )2 and Ni ( ClO 4 )2 Solutions at 25.o0 o •••••••••••••••• • • • • 66 20. Calculated For:ma tion J~1 unction Curves of Co ( 0104 )2 and Ni ( 0104)2 Aquo complexes. • • • • • • • • • • 71 Heat of Reaction of Water with Co(c104 )2 in 1- Butanol at 25°0 ••••••••••.••.. 72 22. Possible Structures of M(C104 )2 in 1-Butanol • • • 75 23. Formation of the .n : 2 Aquocomplex •••••• • • 78 vii CHAPTER I INTRODUCTION Background to the Problem= 1rhe bulk of the scientific research concerned with solution chemistry has been done with aqueous media. However.9 nonaqueous solvent systems often are available which provide better physical and chemical proper ties than does water for particular investigations. In ana= lytical chemistry, for example)} very weak acids or bases are often titrated in a nonaqueous solvent., lvluch nonaqueous polarography :i.s now done 9 and nonaqueous solvents are vitally important to spectroscopy. The technique of solvent extra.a= tion would be impossible without the use of at least one nonaqueous solvento The present research has been undertaken in an attempt ·to learn 1nore of the solution behavior of transition metal salts in a non aqueous rnedi um containing smal 1 am.oun ts of water~ e., g .. 9 cobalt (II) and nickel {II) pe:.rc.h.lora tes in hydrous ]_.,..., butanol, Some studies he:ve indicated that solvat,ion of tran"" sl tion metal perchl.orates in the ii!.lcohol phases after e:x.trac.= tion is prirnar:l..ly due to hydration (28)" The coextra.c·t!:id water in such ca.s es e.xce eds the normB.l coo.rdina tion nu.inbers of the cations, leading to large "hydration nu1~ers". l= Butanol was chosen as the solvent for the study because it is the alcohol of lowest molecular weight which can be used 1 2 as an extractant of aqueous solutions 9 and it also has very favorable thermal properties 9 which will be discussed latero Quite a variety of opinion exists as to the detailed nature of transition metal aquocomplexes in nonaqueous sol= vents (89 151 17 9 19)0 One of the best approaches to a theoret= ical treatment of such species9 and without which no argument would be really cornplete 9 is through thermodynamic informa= tiono Formation constants have been determined for a very limited number of transition metal aquocomplexes in only a few solvents and solvent mixtures (8 9 15,179 19 9 20,21)0 The method of Bjerrum (2) has been most widely used for calcu= lating the formation constantso Spectroscopy is the method most often chosen for aquo= complexing studieso However.,, spectroscopic information is inherently linked to molar extinction coefficients which nearly always are small in the visible region for aquocom= plexeso It is generally difficult to find a metal=water= solvent system which has an infrared absorption band depen= dent only on the wa.ter=metal interactions and not co.arplicated by other vibratlonal modeso At best the data has consider= able scatter which leads to ·uncertai.nty in the formation constants (8yl7)o Polarography has been used as a companion rnethod to spectroscopyj) but some reser"Vations a.lso must be held about its use (17s,20)o An inherent difficulty in using pol.arogra.phy for studying aquo complexing is that a high oonoen t:ration of a. second or supporting electrolyte is necessary. Whe:reas su.c:h an electrolyte rnight not strictly affect the po1a.rographic 3 measurements as such.? it is difficult to imagine that the chemical hydration process would be totally unaffected. Nel= son and Iwamoto {20) have recognized that lithium perchlorate might compete with copper (II) perchlorate for free water in the solution, but stated that no quantitative evidence was available to show that it did; they9 therefore 9 ignored the (likely) interaction of lithium perchlorate in the interpre= tation of their polarographic results. Even at best, formation constants 9 with perhaps a temper ature dependence., can be determined by spectra scopy and polar= ography. However, these two methods give no information about the stepwise enthalpies and entropies of formation, which usually give the best insight in to the nature of the species. To obtain a complete ther.modyna.mic description~ it is necessary to use some form of calorimetry. In order for Bj errum1 s method to be used for obtaining successive fo:r>me.tion constants)) the system should be capable of passing through successive equilibrium states as repetitive additions a:r.e made to the same solution. Thermometric ti tra tions off er a convenient linking of these two requirements. Nio st analytical procedures are ultimately free energy rnethods 9 i.e. 9 they depend on a. property related to the equilibrium constants. Ii'or exa1rrple.9 the pH change in a convention al acid-base titration depends on the relevant ionization constants of the systemq But reaction systems with a sma.11 .free energy change are al so oapabl e of measure= ment if the entropy term in equation (1.1) is favorable. AH= ~G + TAS (lol) 4 Since a measurable heat of reaction is a very general property of chemical processesp thermometric titrations have considerable applicability to both ther.modynamic studies and analytical chemistry.. One of the advantages of using thermo metric titrations

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