Open Collections

UBC Theses and Dissertations

UBC Theses Logo

UBC Theses and Dissertations

Kinetics of some oxidation-reduction reactions in aqueous solutions. Harkness, Alan Chisholm 1963

Your browser doesn't seem to have a PDF viewer, please download the PDF to view this item.

Item Metadata

Download

Media
831-UBC_1963_A1 H2 K4.pdf [ 5.83MB ]
Metadata
JSON: 831-1.0062052.json
JSON-LD: 831-1.0062052-ld.json
RDF/XML (Pretty): 831-1.0062052-rdf.xml
RDF/JSON: 831-1.0062052-rdf.json
Turtle: 831-1.0062052-turtle.txt
N-Triples: 831-1.0062052-rdf-ntriples.txt
Original Record: 831-1.0062052-source.json
Full Text
831-1.0062052-fulltext.txt
Citation
831-1.0062052.ris

Full Text

KINETICS OF SOME OXIDATION-REDUCTION REACTIONS IN AQUEOUS SOLUTIONS by ALAN CHISHOLM HARKNESS B.A., Un i v e r s i t y of B r i t i s h Columbia, 1947 M.Sc, Univ e r s i t y of B r i t i s h Columbia, 1957 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in the Department of CHEMISTRY We accept t h i s thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA March, 1963 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make i t freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of Chemistry  The University of British Columbia, Vancouver 3, Canada. Date March 28, 1963. PUBLICATIONS K i n e t i c s of the Oxidation of Uranium (IV) by Thallium ( I I I ) . , J.Am.Chem.Soc., 81, 3526 (1959). A.C. Harkness and J . Halpern. Medium E f f e c t s i n the Homogeneous C a t a l y t i c A c t i -v a t i o n of Molecular Hydrogen by Metal S a l t s . I I I . S i l v e r and Mercuric S a l t s . ; J.Am.Chem.Soc. , 8_1_, 5854 (1959),. A.J. Chalk, J . Halpern and A.C. Harkness. Spectra of Some T r a n s i t i o n Metal Ions- and Complexes i n D20., J.Chem.Phys., 31, 1147 (1959) J. Halpern and A.C. Harkness. Oxidation of Carbon Monoxide by Metal Ions, J.Am.Chem.Soc., 83, 1258(1961). A.C. Harkness and J . Halpern. The Un i v e r s i t y of B r i t i s h Columbia FACULTY OF GRADUATE STUDIES PROGRAMME OF.THE FINAL ORAL EXAMINATION FOR THE DEGREE OF DOCTOR OF PHILOSOPHY ' ' of ALAN CHISHOLM HARKNESS B.A., The Un i v e r s i t y of B r i t i s h Columbia, 1947 M.Sc, The Uni v e r s i t y of B r i t i s h Columbia, 1957 TUESDAY. APRIL 30th, 1963, AT 9:30 A.M. IN ROOM 261, CHEMISTRY BUILDING COMMITTEE IN CHARGE Chairman: F.H. Soward W.A. Bryce C.A. McDowell B.A. Dunell . E; Peters L.G. Harrison R. Stewart External Examiner:. H. Taube Stanford U n i v e r s i t y KINETICS OF SOME OXIDATION-REDUCTION REACTIONS IN AQUEOUS SOLUTIONS ABSTRACT The k i n e t i c s of the. e l e c t r o n t r a n s f e r r e a c t i o n U(IV) + T l ( I I I ) i U(VI) + T1(I) were examined i n aqueous p e r c h l o r i c a c i d s o l u t i o n . The rate law was found to be of the form -d/u(iv)7 = ^4+7 m3fj j^ZSt?- 1 +k./H+7_2j The two rate constants were i d e n t i f i e d with r e a c t i o n paths involving activated,complexes of the compositions (U'0H*T1) 6 + and (U«0"T1)' 5 +, r e s p e c t i v e l y . The corres-ponding heats and entropies of a c t i v a t i o n , evaluated from rate measurements over the temperature range 16 to 25°, a r e ^ H i # =24.6 kcal/mole, A H 2 # =21.7 kcal/mole, AS]# = 16 e.u. and A S 2* = 7 e.u. The e f f e c t of i o n i c strength and the s p e c i f i c e f f e c t s of various anions and cations on the rate were examined. The r e s u l t s suggest, but do not prove, that the r e a c t i o n occurs through a s i n g l e two-equi-valent step rather than through successive one e l e c t r o n changes. The homogeneous oxidation of carbon monoxide by metal ions i n aqueous solutions was studied. At temperatures below 80° only Hg 2+ and M11O4- were found to oxidize carbon monoxide. The ions Cu 2+, Ag+, Hgo 2 +, Fe 3+, T l 3 + and Cr 20y were i n a c t i v e . K i n e t i c measurements of the r e a c t i o n 2 H g 2 + + CO + H 20 — — > H g 2 2 + + C0 2 + 2H+ i n d i l u t e p e r c h l o r i c a c i d over the temperature range 26 to 54° yielded the rate law zAML = kZco7 /Hg247 The a c t i v a t i o n parameters are AH' f = 14.6 kcal/mole and A S ^ = -13 e.u. It i s be l i e v e d that the r e a c t i o n proceeds by a mechanism which involves the i n s e r t i o n of CO between Hg"*" and a coordinated water molecule, -Hg ~0H 2 2J + CO ^—>. -Hg -c - 0H+ + H + -Hg - C - 0H + > Hg + C0 2 + H+ ( f a s t ) Hg + H g 2 + > Hg 2 2+ ( f a s t ) The oxidation of carbon monoxide by Mn04~ was found to proceed r e a d i l y over the temperature range 28 to 50°. The rate law was found to be = k/TOj /TMn047 with A H * = 13 kcal/mole and A S# = -17 e.u., both sub-s t a n t i a l l y constant over the pH range 1 to 13. The rate determining step i s considered to be the formation of hypomanganate Mn04" + CO + H 20 > Mn0 4 3 _ + C0 2 + 2H+ . which then undergoes further f a s t reactions to y i e l d MnO^2" i n basic s o l u t i o n and Mn0 2 i n a c i d and n e u t r a l s o l u t i o n s . A remarkable feat ure of the l a t t e r r e a c t i o n i s i t s very marked s e n s i t i v i t y to c a t a l y s i s by Ag+ and Hg 2+ (but not by Cu 2+, Cd 2+, Fe 3+, or T l 3 + ) . The rate law of the catalyzed path i s , i n each case, -d/co7 = k /- c o 7 /gno4:7 /K/ where M = A g + or Hg2+. For Ag+, k at 0° i s 1.10 x l O 5 M"2 s e c " 1 with A H * = 1.3 kcal/mole and As# = -30 e.u. For Hg 2 +, k at 0° i s 1.09 x 10 3 M~2 s e c " 1 with AH*= 6.5 k c a l / mole and AS* = -21 e.u. I t i s suggested that the remark-ably high r e a c t i v i t i e s exhibited by carbon monoxide i n these c a t a l y t i c reactions are r e l a t e d to favourable o x i -dation paths involving intermediates such as 0 -Ag - C ••- 0Mn03. GRADUATE STUDIES F i e l d of Study: Physical Inorganic Chemistry Advanced inorganic chemistry H.C Clark Molecular structure C. Reid, R.. Hochstrasser S t a t i s t i c a l mechanics L.G. Harrison Surface chemistry J. Halpern, L.G. Harrison . Chemical k i n e t i c s W.A.Bryce, G.B. Porter, J.. Halpern Ph y s i c a l organic chemistry ; R. Stewart Related studies: : Theory and applications of d i f f e r e n t i a l equations C.A. Swanson Atomic physics M. Bloom Theory of measurements J . Prescott Metal physics J.A.H. Lund ( i ) ABSTRACT The k i n e t i c s of the electron t r a n s f e r reaction U(IV) + T1(III) > U(VI) + T1(I) were examined i n aqueous p e r c h l o r i c acid s o l u t i o n . The rate law was found to be of the form - d - • •- 1 - 2 1 The two rate constants were i d e n t i f i e d with reaction paths i n v o l v i n g a c t i -vated complexes of the compositions (TJ«0H»T1) and (U'OTl) , r e s p e c t i v e l y . The corresponding heats and entropies of a c t i v a t i o n , evaluated from rate measurements over the temperature range 16 to 25°, are AH]* = 24.. 6 kcal/mole, /\Krf ~ 2 1 » 7 kcal/mole, = 16 e.u. and^S* = 7 e.u. The e f f e c t of i o n i c strength and the s p e c i f i c e f f e c t s of various anions and cations on the rate were examined. The r e s u l t s suggest, but do not prove, that the reaction occurs through a s i n g l e two-equivalent step rather than through successive one electron changes. The homogeneous oxidation of carbon monoxide by metal ions i n aqueous solutions was studied. At temperatures below 80° only Hg^ + and MnO^~ were 2H" *4" 2*^~ 3"t" 3"^~ found to oxidize carbon monoxide. The ions Cu , Ag , Hg 2 , Fe , T l and 2-CrgOr, were i n a c t i v e . K i n e t i c measurements of the re a c t i o n 2 H g 2 + + CO + ILjO > H g 2 2 + + C0 2 + 2H + ( i i ) i n d i l u t e p e r c h l o r i c aci d over the temperature range 26 to 54° yie l d e d the rate law -*m =k[C0][Hg2 +J The a c t i v a t i o n parameters are A H = 14..6 kcal/mole and A S =-13 e.u. I t i s believed that the reaction proceeds by a mechanism which involves the i n -s e r t i o n of CO between Hg 2* and a coordinated water molecule, 0 II 2+ k . + + -Hg - 0H 2 + CO > -Hg - C - OH + H 0 II -Hg - C - 0H + ) Hg + C0 2 + H + ( f a s t ) Hg + H g 2 + > H g 2 2 + ( f a s t ) The oxidation of carbon monoxide by MnO^ was found to proceed r e a d i l y over the temperature range 28 to 50°. The rate law was found to be dt ^ J L 4 with A H * = 13 kcal/mole and A S * = -17 e.u., both s u b s t a n t i a l l y constant over the pH range 1 to 13. The rate determining step i s considered to be the formation of hypomanganate MnO^" + CO + H 20 > MnO^3" + C0 2 + 2H + 2-which then undergoes f u r t h e r f a s t reactions to y i e l d MnO^ i n basic s o l u t i o n and Mn02 i n acid and neutral solutions, t A remarkable feature of the l a t t e r r e a c t i o n i s i t s very marked s e n s i t i v i t y to c a t a l y s i s by A g + and H g 2 + (but not by C u 2 + , C d 2 + , Fe^ +, or ( i i i ) 3+ T ] / ). The rate law of the catalyzed path i s , i n each case, -dgQi = k [ c o ] r M n 0 ^ j ( ; M ] where M = A g + or Hg2"1". For Ag +, k at 0° i s 1.10 x 10 5 M~2 s e c " 1 with A B * = 1.3 kcal/mole and A S* = -30 e.u. For Hg2"*", k at 0° i s 1.09 x 10 3 M"*2 s e c " 1 with A H * = 6.5 kcal/mole and A S * = -21 e.u. I t i s suggested that the remarkably high r e a c t i v i t i e s exhibited by carbon monoxide i n these c a t a l y t i c reactions are r e l a t e d to favourable oxidation paths i n v o l v i n g intermediates such as 0 m -Ag - C - OMnO,. (x) ACKNOWLEDGEMENTS I wish to express my appreciation f o r the assistance given me by Dr. J . Halpern i n a l l aspects during the course of t h i s work. The awards of a Fellowship from the Consolidated Mining and Smelting Company of Canada Limited and of a Sbholarship from the B r i t i s h Columbia Sugar Refining Company were most h e l p f u l and are greatly ap-preciated. Above a l l a great debt i s owed to my family. ( i v ) TABLE OF CONTENTS Page No. KINETICS OF SOME OXIDATION-REDUCTION REACTIONS IN AQUEOUS SOLUTIONS 1 GENERAL INTRODUCTION 1 PART I -THE KINETICS OF THE OXIDATION OF URANIUM (IV) BY THALLIUM ( I I I ) 12 Introduction 12 Experimental 26 Materials 26 A n a l y t i c a l 28 K i n e t i c a l Measurements 29 Results and Discussions 30 PART I I -KINETICS OF THE HOMOGENEOUS OXIDATION OF CARBON MONOXIDE BY METAL IONS 62 Introduction 62 Experimental 67 Materials 67 A n a l y t i c a l 68 Procedure 68 Results and Discussion 74 Oxidation of CO by H g ( l l ) 74 Oxidation of CO by MnO^" 85 C a t a l y s i s of the CO - MnO^~ Reaction by Ag(l) and H g ( l l ) 94 Tracer Studies with 0 1 8 109 Mechanism of the CO - H g ( l l ) Reaction 113 Results of Related CO Reactions Mechanisms of the Uncatalyzed and Catalyzed CO - MnO^ Reactions Comparison of CO, Hg and HCOOH as Reductants Comparison of the I s o e l e c t r o n i c Species CO, N_ and CN ( v i ) LIST OF TABLES Table Page No. No. I Exchange rate constants for thallium chloro complexes 21 II Stoichiometry of U(IV) - Tl(lII) reaction 32 III Effect of in i t i a l concentrations of U(IV) and Tl(III) on second order rate constant of U(IV) - T l ( l l l ) reaction 34-IV Effect of NaCl on rate of U(IV) - Tl(lII) reaction 35 V Effect of Na SO on rate of TJ(lV) - Tl(III) reaction 38 2 4 VI Effect of some metal ions on rate of U(IV) - Tl(IIl) reaction 40 VII Effect of NaClO^ on rate of U(IV) - Tl(III) reaction 41 VIII Effect of HCIO^ on rate of U(IV) - T l ( l l l ) reaction 44 IX Effect of HCIO^ on rate of U(IV) - Tl(lII) reaction 45 X Effect of HCIO^ on rate of U(IV) - T l ( l l l ) reaction 46 XI Effect of HCIO^ on rate of U(IV) - Tl(lII) reaction 47 XII Effect of HCIO^ on rate of U(IV) - T l ( l l l ) reaction 48 XIII Hydrolysis constants of and T l 3 + 54 XIV Kinetic data for the oxidation of U(IV) by Tl(lII) 54 XV Variation of rate of U(IV) - T l ( l l l ) reaction with temperature 56 XVI Solubility of CO in water 75 XVII Effect of reactant concentrations on rate of the CO - Hg(Il) reaction 78 XVIII Effect of HCIO^ on the rate of the CO - Hg(ll) reaction 79 XIX Effect of NaClO^ on the rate of the CO - Hg(II) reaction 80 XX Activation parameters of the CO - Hg(ll) reaction at 40.0° 82 XXI Effect of temperature on the CO - Hg(ll) reaction 83 XXII Kinetic data for the oxidation of CO by MnO^ " in acid and neutral solutions at 50,0° 86 ( v i i ) Table Page No. No. XXIII E f f e c t of temperature on the rate of the oxidation of CO by MnO^" i n acid and neutral solutions 87 XXIV K i n e t i c data f o r the CO - MnO^~ reac t i o n i n a l k a l i n e s o l u t i o n 93 XXV K i n e t i c parameters f o r the CO - MhO^" reac t i o n at 50.0° 94 XXVI K i n e t i c data f o r the H g ( l l ) catalyzed oxidation of CO by MnO." 102 4 XXVII K i n e t i c data f o r the Ag(I) catalyzed oxidation of CO by MnO^" 103 XXVIII E f f e c t of temperature on the rate of the H g ( l l ) catalyzed CO - MhO^" reac t i o n 104 XXIX E f f e c t of temperature on the rate of the Ag(l) catalyzed CO - MnO^" re a c t i o n 105 XXX A c t i v a t i o n parameters f o r the catalyzed CO - Mn0^~ reactions 108 18 XXXI Summary of 0 t r a n s f e r experiments 114. XXXII Summary of k i n e t i c data f o r some oxidations of H 2, HCOOH, HCOO" and CCT 125 (viii) LIST OF FIGURES' Fig. Page No. No. 1. Profile of potential energy surfaces showing intersection of surfaces corresponding to reactants and products 9 2. Rate of U(IV) - T l ( l l l ) reaction in 3M HCIO^ at 25.0° 33 3. Effect of Cl" on rate of TJ(IV) - Tl(III) reaction 36 4. Effect of SO^2" on rate of U(IV) - T l ( l l l ) reaction 39 5. Effect of NaClO^ on rate of TT(IV) - T l ( l l l ) reaction 42 6. Rate plots of U(IV) - Tl(lII) reaction for various HCIO^ concentrations 49 7. Effect of HCIO^ on second order rate constant of U(IV) - T l ( l l l ) reaction 50 8. Effect of acidity on rate of U(IV) - Tl(IIl) reaction at various temperatures 53 9. Arrhenius plots for rate constants and kg of U(IV) - Tl(lII) reaction 55 10. Gas contacting apparatus 69 11. Gas absorption apparatus 73 12. Rate plots for CO - Hg(ll) reaction at 4-0.0° 77 13. Effect of CO pressure on rate of CO - Hg(ll) reaction 81 14. Arrhenius plots for CO - Hg(II) reaction 84 15. Rate plots for CO - MnO^ " reaction in acid and neutral solutions at 50.,0° 88 16. Effect of CO pressure on rate of CO - MnO^ ~ reaction 89 17. Rate of CO absorption by alkaline MnO^ ~ 91 18. Rate plot for data of Figure 17 92 19. Arrhenius plot for CO - MnO," reaction in acid solution 95 (ix) F i g . Page No. No. 20. Arrhenius p l o t f o r CO - MnO^~ reaction i n neutral s o l u t i o n 96 21. Arrhenius p l o t f o r CO - MnO^ re a c t i o n i n a l k a l i n e s o l u t i o n 97 22. Rate plo t s f o r uncatalyzed and catalyzed CO - MhO^~ reaction at 13,0° 100 23. Arrhenius p l o t f o r H g ( l l ) catalyzed CO - MnO ~ reaction 106 4 24., Arrhenius p l o t f o r Ag(l) catalyzed CO - MnO 4" reaction 107 25, Apparatus f o r decomposing BaCO, and c o l l e c t i n g CO- 110 1. KINETICS OF SOME OXIDATION-REDUCTION REACTIONS IN AQUEOUS SOLUTIONS General Introduction This t h e s i s i s concerned with the k i n e t i c s of some oxidation-reduction reactions i n aqueous solutions and comprises two parts. Part I describes an i n v e s t i g a t i o n of the k i n e t i c s of the reduction of thallium (III ) by uranium (IV), This i s a reac t i o n between a two equivalent oxidant and a two equivalent reductant. One of the objects of the i n v e s t i g a t i o n was to determine whether i t proceeds i n a si n g l e step or i n successive one ele c t r o n steps. Part II describes a k i n e t i c i n v e s t i g a t i o n of the reduction of some metal ions by the molecular reducing agent, carbon monoxide. This was p r i m a r i l y an exploratory i n v e s t i g a t i o n o f ' a new clas s of reactions r e l a t e d i n some respects to the homogeneous reactions of molecular hydrogen which have been extensively investigated i n t h i s laboratory ( l ) . In recent years great i n t e r e s t has developed i n oxidation-reduction or e l e c t r o n t r a n s f e r reactions. The subject has been reviewed by Zwolinski, R. J . Marcus and Eyring (2), Basolo and Pearson (3), Taube (4)* George and G r i f f i t h (5), Stranks (6)" and Halpern (7). Oh the basis of experimental evidence i t has been possible to d i s -t i n g u i s h two f a i r l y d e f i n i t e forms of the activated complex f o r e l e c t r o n t r a n s f e r reactions between metal ions. These are defined by the r e l a t i v e configuration of the ligands i n the complex. The two types represent extremes i n behaviour and i n only a comparatively few cases has i t been possible to assign reactions to either class. The outer sphere complex is characterized by the constancy of the fi r s t coordination spheres of the reactants. That is, both the number and identity of the ligands around each metal ion remain unchanged in the comr plex. This is the mode by which substitution inert pairs react,, Examples are MnO^ " - MnC^2", Fe(CN)^" - Fe(CN)63" and Fe(phen) 3 2 + - Fe(phen) 3 3 +, (Note). The outer sphere complex has also been established for pairs of reactants in which one member is substitution labile. In such a case i t is necessary to establish that the rate law corresponds to an activated com-plex containing the same number of ligands as are held by the metal ions separately. Examples are Co(en)^ - Co(en)<j3 and C r 2 + - CoCNH^)^3 .. The question of how close the reacting pairs approach in this type of complex remains unanswered. There is no direct evidence on this point. Sheppard and Wahl (8) found that substituting CsOH for NaOH as the medium in the MnO^  - MnO^  exchange resulted in a threefold increase in rate. The reason for this is not clear, but i t does suggest that consider-2-ation be given to a cation bridged structure, e.g. oyfnO — Cs — OMnO^  ., The inner-sphere activated complex is characterized by the sharing of ligands by the reacting ions. Its formation is quite analogous to that of ordinary ligand substitution reactions. In this case one of the reactants with its own inner coordination sphere acts as a ligand being (Note)s Abbreviations used for ligands are* phen = 1,10-phenanthroline; en = ethylenediamine; bip = 2,2 -bipyridine 3. substituted into the coordination sphere of the other reactant. In order to demonstrate the existence of this type of complex the coordination spheres of one metal ion reactant and the other metal ion product must be substitu-tion inert. The most extensive investigations of this type of mechanism have been done by Taube and his school on the reduction of Co(ill) complexes by 2+ Cr . Co (III)) and Cr(IIl) complexes are relatively inert to substitution 2+ whereas Cr is very labile. One of the f i r s t demonstrations of an inner-sphere complex was in the reduction of (NH^CoCl2"1" by Cr2"1" (9). The oxidized 2+ 3+ product was found to be CrCl rather than CrCKjO)^ .. This was taken to mean that a Cr - Cl bond was formed in the active complex leading to the structure (NH3)5Co - Cl - CrCEjOj^4" Further evidence for this structure was adduced by repeating the experiment in the presence of radioactive chloride ions {10%. Substantially no radio-2+ activity was found in the product CrCl „, Recently Haim and Wilmarth (11) have reported the isolation of a complex ion resulting from the oxidation of the pentacyano complex of Co(II) 3-by Fe(CN)^ . They considered the product to have the bridged structure (CN)5Fe - CN - Co(CN)5 It must not be assumed that the bridged activated complex always leads to atom transfer as the redox mechanism. For example, in the case of p i p 3+ the reaction between Cr and IrCl^ (10): the products are C r ^ O ) ^ and IrCl^ . The mechanism is considered to involve an inner-sphere complex and spectral evidence suggests this. The point here is that the bridging ligand w i l l follow; that partner which has the greatest a f f i n i t y f o r i t . I t i s apparent that f a i r l y r e s t r i c t i v e conditions must be met-before a r e a c t i o n can be assigned to eit h e r of the foregoing classes with any degree of c e r t a i n t y . Unfortunately reactions of aquo ions i n general do not meet these conditions. The o v e r r i d i n g d i f f i c u l t y i n the case of aquo ions i s i n attempting to define the reacting species. Inferences may be drawn by observing the e f f e c t s of such v a r i a b l e s as a c i d i t y , i o n i c strength and anions, but i t i s r a r e l y possible to do more than t h i s . Some of these points may be i l l u s t r a t e d by considering reactions 2+ 3+ such as the Fe - Fe exchange. Like most electron t r a n s f e r reactions i t i s f i r s t order i n each reactant (12). There i s an inverse f i r s t order dependence on hydrogen ion concentration. This i s quite common and i s us u a l l y taken to i n d i c a t e a rea c t i o n path i n v o l v i n g an hydrolyzed species? i n t h i s case FeOH 2 4. However, t h i s cannot be confirmed as there i s no k i n e t i c d i f f e r e n c e between FeOH 2 + - F e 2 + , F e 3 + - FeOH +, F e 3 + - F e 2 + - GH° and other formulations. The f a c t that the f e r r i c i o n i s more strongly hydro-lyzed than the ferrous i on i s the excuse f o r considering FeOH^ + as a reactant.. A' s i m i l a r s i t u a t i o n p r e v a i l s with regard to the e f f e c t of ch l o r i d e i o n and other anions. The c a t a l y s i s of the reac t i o n (a f a c t o r of 1000) by 0H~ l e d to the suggestion (13) that a hydrogen bridged intermediate may be involved, (H 20) 5Fe0 H - - - OTe(E20j)5*+ H H S i m i l a r l y the intermediate f o r the uncatalyzed reaction i s 5. H (HjOj-Pe 0 - - - H - - - 0 F e ( l L O ) / H H ^ 5 The reactions are completed by t r a n s f e r r i n g the hydrogen atom. Hudis and Dodson found that the rates f o r both paths were lowered by a f a c t o r of 2 when the reaction was conducted i n heavy water. This gives some support to the proposals. However, hydrogen isotope e f f e c t s i n aqueous solutions 2+ are not at present of much diagnostic value. In the reduction of (NH^)^CoCl by Cr which proceeds by a c h l o r i d e bridge there i s a 30% reduction i n rate i n changing the solvent from HgO to D^O (15). The uncertainty of arguing by analogy i s i l l u s t r a t e d i n Taube's attempt (4) to assign a bridged mechanism to the c h l o r i d e catalyzed Fe F e ^ + exchange. Both C r 2 * and F e 2 + reduce several f e r r i c complexes at the same r e l a t i v e rates and chloride catalyzes both the C r 2 * - F e ^ + and F e 2 + — Fe-7 reactions. CrCl i s formed q u a n t i t a t i v e l y i n the one case (10), but the corresponding F e C l 2 + i s l a b i l e . The C r 2 + reduction of C r ^ O ) ^ " 4 " i s not catalyzed by c h l o r i d e and CrCl i s not formed (16). This evidence was i n t e r -preted to mean that chloride must act as a bridging agent to e f f e c t c a t a l y s i s 2+ 3+ and, therefore, the Fe - Fe^ r e a c t i o n has a bridged activated complex. I t has since been shown (16A) that the Cr2"1" - F e ^ + r e a c t i o n takes place without the p r i o r formation of F e C l 2 + . Also the oxidation of V 2 + by s u b s t i t u t i o n i n e r t C o ( l l l ) complexes i s catalyzed by chloride as i s the reaction between Cr 2 4" and Co(NH^)^ (17). These r e s u l t s show that chloride does not have to act as a bridging l i g a n d to exert a c a t a l y t i c e f f e c t and the previous conclusions concerning the Cr2"4" - Fe-* + and F e 2 + - F e ^ + reactions are v i t i a t e d . I n c i d e n t a l l y the experiments with showed that i t s reactions 6. were c l o s e r to those of CrCbip^ 2" 1" than those of C r 2 + , The i m p l i c a t i o n i s that V2"*" reacts through outer-sphere complexes j at l e a s t with s u b s t i t -u t i o n i n e r t o x i d i z i n g agents. I t i s of course not necessary that a bridging ligand be an anion. 18 The 0 t r a c e r work of Murmann, Taube and Posey (16) shows that a substantial 3+ 2"^" portion of the water i s t r a n s f e r r e d i n the reduction of (NH^J^CoHgO by Cr . 3+ I f (NH^jyJoHgO can be considered an aquo ion t h i s i s the only case where the mechanism of a r e a c t i o n between aquo ions i s known. In many cases the c a t a l y t i c e f f e c t s of anions are at t r i b u t e d to a possible bridging mechanism. While t h i s may be so f o r a p a r t i c u l a r reaction or series of reactions i t i s by no means u n i v e r s a l . For example, the redue-.3+ w nJ2+ t i o n of (NH^J^CoOHg by Cr i s accelerated by pyrophosphate which appears 2+ i n the C r ( l l l ) product (18). When the corresponding (NH^J^CoCl i s reduced both c h l o r i d e and pyrophosphate are found i n the C r ( l I I ) product. 2+ The reductions of various pentamino Co(III) complexes by Cr(bip)3 proceed by outer-sphere complexes (19). Here there i s marked v a r i a t i o n i n rate brought about jqy d i f f e r e n t anions. This may be due to what are termed medium or s a l t e f f e c t s . Medium e f f e c t s are r e l a t e d to such factors as i o n i c strength and the nature of the supporting e l e c t r o l y t e . The influence may be exerted through changes i n a c t i v i t y c o e f f i c i e n t s or ease of adjustment of i o n i c atmos-pheres i n forming the activated complex. This aspect of k i n e t i c s i n general i s not well understood and can lead to d i f f i c u l t i e s of i n t e r p r e t a t i o n . The data of Bonner and Hunt (20) on the C o 2 + - C o 3 + exchange show that the rate decreases as HCIO^ i s increased at constant i o n i c strength. The r e s u l t can be int e r p r e t e d equally w e l l as e i t h e r a medium e f f e c t or an hydroxyl path. 7. It is useful to consider the theoretical aspects of electron transfer insofar as light is shed on the importance of various factors. One of the earliest advances was made by Libby (21) who suggested that the Franck-Condon principle is applicable to electron transfer processes in solution. This principle states simply that as electronic motion is much faster than atomic motion an electronic transition leaves the atomic con-figuration unchanged. This has led to the recognition that atomic reorgani-sation provides a barrier to electron transfer. Some controversy has arisen as to whether or not electron transfer precedes or follows atomic reorgani-zation. It would appear that the first law of thermodynamics requires that substantial reorganization take place prior to electron transfer. Otherwise the electron transfer would result in the production of two energy rich species (an oxidized ion in the environment of a reduced ion and vice versa) which would give off energy in attaining their equilibrium configurations. Energy is conserved i f both reactants assume an intermediate configuration prior to electron transfer. The energy released to the medium as the products assume their equilibrium configuration is balanced by the energy taken up by the reactants. The Franck-Condon restriction is most pronounced for reactions in which there is no free energy change. R. J. Marcus, Zwolinski and Eyring (22, 2) developed a treatment of electron transfer based on a tunnelling model. In this treatment the free energy of activation is broken down into three components % an electrostatic term representing the free energy of repulsion between the reacting ions, the free energy of reorganization of the ionic environments and the probab-i l i t y of the electron penetrating the coulombic barrier. The distance of approach of the reactants is chosen to effect a compromise between the repulsion term and the probability term. No attempt was made to calculate 8. 2+ the reorganizations! term. I t was taken as constant by choosing the Fe 3+ - Fe exchange as a standard and making the calculated and experimental values of the f r e e energy of a c t i v a t i o n agree. The quantitative r e s u l t s obtained on t h i s model are of l i m i t e d s i g n i f i c a n c e . Being made to agree with one value may be the reason f o r near agreement with others, but t h i s may be f o r t u i t o u s . Thus the calculated value f o r the f r e e energy of a c t i v a t i o n of + 2+ the NpOg - NpCv, exchange was I 4 .6 kcal/mole which i s i n ex c e l l e n t agree-ment with the experimental value of 14.3 kcal/mole (23). However, Cohen et a l (24) repeated the exchange r e a c t i o n while varying the d i e l e c t r i c constant from 68 to 88. Contrary to the predictions of the theory, they found that the rate of exchange remained constant. R. A. Marcus i n a serie s of papers (25, 26, 27, 28) has attempted to formulate a quan t i t a t i v e theory taking i n t o account a l l the i n t e r a c t i o n s of the reactants and the medium. In general, a system can be represented by a po t e n t i a l energy surface i n many atomic coordinates. A reaction can be represented by the i n t e r s e c t i o n of two surfaces corresponding to the system of reactants and that of products. A p r o f i l e of the two surfaces and t h e i r i n t e r s e c t i o n i s shown i n Figure 1. There are four casesi 1. No e l e c t r o n i c i n t e r a c t i o n . A system w i l l stay on i t s own surface on passing through the i n t e r s e c t i o n and no reaction takes place. 2. Strong e l e c t r o n i c i n t e r a c t i o n . This leads to sub s t a n t i a l s p l i t t i n g of the surfaces and appreciable lowering of the a c t i v a t i o n energy which i s at present not c a l c u l a b l e . 3. Weak i n t e r a c t i o n , adiabatic transfer. S p l i t t i n g of the surfaces occurs so that a system passing through the i n t e r -section w i l l remain on the lower surface. The i n t e r a c t i o n energy i s small 9. Products Atomic configuration Figure 1. P r o f i l e of p o t e n t i a l energy surfaces showing i n t e r s e c t i o n of surfaces corresponding to reactants and products. 10. enough to be ignored in the calculations. 4. Very weak interaction, non-adiabatic transfer. This is like the previous case except that there is now a probability that the system will jump from the lower to the upper surface on passing through the intersection. The free energy of activation is, as by Zwolinski et al, split into components. AT* = A F j + AF^ - RTlnk Here /SF-^ ^ s the free energy of repulsion of the reactants and^AFg isi the reorganizational free energy, k is the probability that the system will remain on the lower surface. In the adiabatic case k = 1. The complete evaluation of these quantities is not easy, but fairly simple expressions are obtained by making these approximations. Each reac-tant including its coordination shell is treated as a rigid sphere of radius r. The separation of the reactants in the activated complex is the sum of their radii (a = r^ + r 2 ) . The results are 1 Da and AFg* = m 2 E where t 2 r l 2 y2 ffj(n2 "'j e^, e 2 are the charges of the reactants and e-^ ', e 2' are those of the products, Ae represents the number of electrons transferred, D is the dielectric con-stant and n is the refractive index of the solvent. 11. For an exchange reaction the free energy of a c t i v a t i o n i s given by 1~ 2r F a i r agreement i s obtained between experimental and c a l c u l a t e d •» 2— values f o r reactions such as the MnO^° - MnO. exchange i n which the atomic configurations of the i n i t i a l and f i n a l states are nearly i d e n t i c a l . In the case of aquo ions the ebove expressions are not applicable because ligand motion has not been considered. The more complete theory i s supposed to account f o r t h i s . rather than to the inner-sphere model of the activated complex, because of the r e s t r i c t i o n on a low i n t e r a c t i o n energy. In as much as free energies of repulsion and reorganization are involved i n both models, however, the q u a l i t a t i v e conclusions are common. Perhaps the most serious omission of the theory i s i t s f a i l u r e to take i n t o account the i n d i v i d u a l i t y of the reactants and ligands. These are probably r e f l e c t e d i n the ^ F ° values, but t h i s term vanishes i n exchange reactions. Recently Hush (29) has presented a theory which i s quite s i m i l a r to that of Marcus. In t h i s theory the parameter, m, which Marcus introduced as a Lagrangian m u l t i p l i e r , takes on the physical s i g n i f i c a n c e of an electron p r o b a b i l i t y density. The theory corresponds more c l o s e l y to the outer-sphere model 12, PART I THE KINETICS OF THE OXIDATION OF URANIUM (IV) BY THALLIUM ( I I I ) Introduction In many redox reactions at l e a s t one of the reactants changes i t s oxidation state by more than one u n i t . The r e s o l u t i o n of the o v e r a l l mechanism in t o elementary steps makes i t possible to determine how many electrons are transferred i n a s i n g l e step. The question of one versus two el e c t r o n t r a n s f e r s has been d i s -cussed f o r some time. K i r k and Browne (30) f i r s t c l a s s i f i e d a number of inorganic o x i d i z i n g agents according to the stoichiometries of t h e i r reac-t i o n s with hydrazine, Higginson, Sutton and Wright (31, 32, 33) c l a r i f i e d and expanded t h i s c l a s s i f i c a t i o n , Higginson and Marshall (34) showed that the products of the oxidation of sulfurous acid also depended on the nature of the o x i d i z i n g agent. This d i s c r i m i n a t i n g action can be represented schematically i n the fol l o w i n g manner, 1. S(n) > S(n+1) 2. 2S(n+l) > |[S(n+l)J 2 3. 2S(n+l) ) S(n+2) + S(n) 4. S(n+1) > S(n+2) 5. S(n) > S(n+2) The substrate, S, e x i s t s i n oxidation state n, A one-equivalent o x i d i z i n g agent produces the r a d i c a l , S(n+l) (Step l ) . This r a d i c a l may then 13. dimerize (Step 2), disproportionate (Step 3) or undergo a f u r t h e r one-equivalent oxidation (Step 4 ) . Thus a one-equivalent o x i d i z i n g agent acting on S(n) w i l l produce two products S(n+l)2 and S(n+2). A two-equivalent o x i d i z i n g agent w i l l produce only one product, S(n+2) (Step 5). With hydrazine the i n i t i a l products are not stable and fu r t h e r reactions follow. The two-equivalent agents produce Ng and the one-equivalent agents produce both N2 and NH^. Two-equivalent N 2 H 4 -— > N 2H 2 + 2 H + ¥ * 2 -— > N 2 + 2 H + One-equivalent . N 2 H 4 —> N 2H 3 + H + (a) N 2H 3 -—> N 2H 2 + H + —> N 2H + H + N 2H —> N 2 + H + (b) 2 N 2H 3 — > N 4 H 6 N 4 H 6 " — • N 2 + 2 N H 3 The oxidation of sulfurous acid produces s u l i Two- equi valent so 3 2~ -— > so3 (so/ -) One-equivalent S 0 3 2 " -— > so3" 2 S 0 3 ~ —> S 2 ° 6 2 " On the basis of the observed stoichiometries three classes of ox i d i z i n g agents were distinguished. One-equivalent% Ce(lV), C o ( l I I ) , F e ( l l l ) , OH Two-equivalent» I 2 , B r 2 , C l 2 , BrO", B r 0 3 " , I 0 3 " , HgO.,, T 1 ( 1 1 1 ) 2 - ? « . -4-Mixed-equivalentt ^T2°7 » M a 04. » P ^ O l ^ , V 0 2 H . From these and other observations Higginson and Marshall put f o r t h these generalizations* 1) Redox reactions between two t r a n s i t i o n metal ions u s u a l l y occur i n one-equivalent steps. 2) Redox reactions between d e r i v a t i v e s of two non-transition elements u s u a l l y occur i n two-equivalent steps. 3) Redox reactions between t r a n s i t i o n metal ions and d e r i v a t i v e s of non-transition elements may occur i n e i t h e r one- or two-equivalent steps with the one-equivalent mechanism predomi-nating. U) I f one of the reactants i s a free r a d i c a l a one-equivalent r e a c t i o n i s more l i k e l y . From a consideration of the c h a r a c t e r i s t i c s of the oxidation reduc-t i o n reactions of organic compounds, p a r t i c u l a r l y quinones, Michaelis (35) proposed that any oxidation has to proceed i n successive univalent steps, i . e . with the formation of free r a d i c a l s . Westheimer (36) has shown that t h i s i s not n e c e s s a r i l y so with e i t h e r organic or inorganic reactions. He also points out that there i s probably no t h e o r e t i c a l grounds f o r objecting to a two el e c t r o n t r a n s f e r between ions i n polar solvents. The major objection to a two electron t r a n s f e r appears to have arisen through the assumption that such a t r a n s f e r i s very improbable quantum mechanically (e.g. 37). I f atom tr a n s f e r i s involved, even i n d i r e c t l y as with solvent molecules, there would appear to be no r e s t r i c t i o n on the number of electrons transferred. In t h i s connection, both theory and experiment put the cross section f o r two electron t r a n s f e r between rare gas atoms and t h e i r ions lower by only a f a c t o r of 2 to 4 than f o r the corresponding s i n g l e e l e c t r o n t r a n s f e r (38). Shaffer compared the r e l a t i v e rates of several redox reactions and 15. was moved to enunciate a principle of equi-valence change (39, 4-0). This states that only those reactions w i l l be fast i n which the oxidation states of the reactants change by an equal number, i.e. complementary reactions. Non-complementary reactions were considered to be slow in the absence of catalysts. For example, Ce(lV) i s reduced to C e ( l l l ) very rapidly by the one-equivalent reducing agents T i ( l I I ) , F e ( l l ) , Br". On the other hand the 2+ actions of the two-equivalent agents T l ( l ) , As(IIl), Hgg , H^ POg B T Q slow. Similarly the reductions of T l ( l l l ) by Sn ( l l ) , Hg 2 2 +, ^ 8 0 3 are fast while those by Fe(Il) and T i ( l l l ) are slow. One implication of this principle i s that reactions between two-equivalent reactants involve the transfer of two electrons i n one step. In effect, the principle of equi-valence change states that non-complementary reactions are slow because they must proceed through a termole-cular mechanism or through the formation of unstable valence states. Four + 2+ simple mechanisms are illustrated below for the oxidation of A to A and the reduction of B 2* to B 1) Termolecular 2A + + B 2 + > 2 A 2 + + B 2) Bimolecular, i n i t i a l one-equivalent step A+ + g2+ > A2+ + B + A + + B + > A 2 + + B 3 ) Bimolecular, i n i t i a l two-equivalent step A + + B 2* > A 3 + + B A + + A 3 + > 2 A 2 + 4) Bimolecular, i n i t i a l disproportionation 2 A + > A + A 2 + A + B 2 + > A 2 + + B 16. The main f a i l i n g of the p r i n c i p l e i s that i t ignores other f a c -tors and i t i s not s u r p r i s i n g that there are many exceptions. Remick (41) 3+ suggested that the f a s t oxidations of Sn(II) by Ce(lV), Fe(phen) 3 and 3_ Mo(CN)g are due to the large d i f f e r e n c e s i n p o t e n t i a l s of the couples. Halpern (42) has given examples and has suggested other f a c t o r s of importance. For example, the rate of oxidation of F e ( I l ) by T l ( l l l ) i s intermediate between the corresponding i s o t o p i c electron exchange reactions. This was a t t r i b u t e d to the more p o s i t i v e entropy of a c t i v a t i o n f o r the mixed reacti o n . Another f a c t o r i s that the mixed reaction has a negative f r e e energy change. Another type of v i o l a t i o n of the p r i n c i p l e i s the reduction of Cu by Hg (43). Reaction by e i t h e r mechanism 2 or 3 would give as the rate determining step C u 2 + + Hg » C u + + H + + H (^H° = 54 Kcal) or C u 2 + + Hg > Cu + 2H + ( A H ° = 66 Kcal) but the a c t i v a t i o n energy i s only 26 Kcal, The postulated mechanism i s C u 2 + + Hg CuH + + H + CuH + + C u 2 + > 2Cu + + H + The determining f a c t o r here i s the s t a b i l i z a t i o n of the intermediate by co-valent bonding. Reactions i n which both partners undergo a two-equivalent change are w e l l known. However, i n only a few cases have the k i n e t i c s supported a two e l e c t r o n t r a n s f e r and t h i s i s accomplished through group t r a n s f e r . The reduction of ClO^" by SO^ - i s f i r s t order i n each reactant and 18 by using G Halperin and Taube {UU) showed that there was complete oxygen tra n s f e r i n the r a t e determining step 17. CIO3 + SO34 > C1CT2 + SD^ Stewart (4.5)' has shown that the rate determining step i n the permanganate oxidation of benzhydrol i s the t r a n s f e r of a hydride ion. ( C 6 H 5 ) 2 CHO" + MnC^" > ( C6 H5% 0 0 + m n 0 ^ " Both U(IV) and T l ( l l l ) are normally two-equivalent reactants and i t was f e l t that t h e i r r e a c t i o n might well involve the simultaneous t r a n s f e r of two electrons. In t h i s context, the remarks of Westheimer (36) are pertinent. He considered that there could be no p r a c t i c a l d i s t i n c t i o n between a simultaneous two electron t r a n s f e r and successive one electron t r a n s f e r s i f the second followed before the "intermediate" could break out of i t s s o l -vent cage; a time of the order of I O - 1 1 second. Several reactions i n v o l v i n g the oxidation of TI(lV) to U(VT) have been studied. These include the i s o t o p i c exchange (4-6, 47) and the reactions with other a c t i n i d e s , Pu(IV) (48), Pu(VT) (49) and Np(VI) (50). The o x i -dations with F e ( l l l ) (51) and Ce(lV) (52) have also been studied. The k i n e t i c s of the reactions with the non m e t a l l i c and multi-equivalent o x i d i z i n g agents 02(53), H 20 2 (54) and CIO3" (55) have also been reported. There i s one feature common to a l l the reactions and which follows from the chemistry of uranium. Tetravalent uranium e x i s t s as U 4 + and hexa-valent as TO2 (56), plus associated water molecules. I t i s to be expected that the k i n e t i c s of the oxidation of w i l l show an inverse dependence on a c i d i t y and t h i s i s found to be so. The U(IV) - u(Vl) exchange has some i n t e r e s t i n g features. Rona (47) investigated t h i s i n d i l u t e HC1 solutions and found that the r e a c t i o n was second order i n U 4 , f i r s t order i n UO-j and inverse f o u r t h order i n H +. 18. She suggested the mechanism UOH3+ + U0 2 2 + + 2^0 j — * 0=U(0H> - 0" - U(0H) 2 3 + + 2H+ 0=0"(OH) - 0 U ( 0 H ) 2 3 + + U0H3+ > products which explains the kinetics. The interesting thing here is the postulated dinuclear intermediate. A" dinuclear U(IV) species would lead to the same result. Newton (A-9) postulated a similar intermediate in the U(IV) - Pu(Vl) reaction although i t is f i r s t order in both reactants. Very recently Newton and Baker (57) have shown that the reduction of U(VI) by Cr(ll) proceeds by way of an intermediate which will reduce other oxidants. Masters and Schwartz (46) investigated the uranium exchange in perchloric acid. They found that in addition to the path described by Rona another path predominated at low TI(lV) levels and higher temperatures. This path is described by the rate law Rate = k [U^ +][U0 2 2 +] CH +] It leads to exchange through the formation of U(V) and involves the same activated complex as does the disproportionation reaction of U(V). The reactions of U(IV) with the other cations follow similar mechanisms which can be illustrated with respect to Fe(III) (neglecting the acid dependence). U(IV) + Fe(lII) U(V) + Fe(ll) U(V) + Fe(lII) u(VI) + F e ( n ) There is no direct evidence for this sequence but the arguments are theses A two electron transfer as the rate determining step is unlikely since Fe(l) 19. and Ce(ll) are unknown. For the reactions involving Pu(Vl) and Np(Vl) the second step would be slow and the kinetics would be altered. Granted the one electron f i r s t step, the disproportionation of U(V) in the second step is unlikely as not being consistent with the known rate and steady state concen-tration. Except for the exchange reaction in dilute hydrochloric acid the foregoing reactions a l l show a rate increase with an increase of ionic strength. They are subject to catalysis by sulfate, but not by chloride ions. (The Fe(lII) reaction was not tested)'. The oxidation of U(lVr)' by 0 2 and H 20 2 both appear to proceed by chain mechanisms although the overall reactions are nearly second order (first order in each reactant). The mechanism proposed by Halpern and Smith (53) for the oxidation with 0 2 is u(iv) + °2 " —> U(V) + H 02 U(T) + °2 " • U(VI) + H02 U(IV) + H02 -—> u(v) + H2°2 u(v) + H02 -—> U(VI) + TJ(IV) + *2°2 ~ > TJ(VI) which accounts for the kinetics and the inhibition by Cl and Ag . However, Gordon and Taube (58) found that there is complete transfer of oxygen from 0 2 according to the equation 2U 4 + + 0 2 + H^O > 2U0 2 2 + + A H + The chain mechanism predicts at most 50$ transfer. 20. The most interesting reaction of thallium is the Tl(l) - Tl(IIl) exchange reaction. It has been studied by many workers who have emphasized different aspects. In common with many reactions there is an acid dependent path so that the rate expression may be written Rate = k Q [ T l + j r T l 3 + ] + ^ [ T l + ] [ T10H2"4"] However, there has been lack of agreement as to which path is the more important. The very careful work of Roig and Dodson ( 5 9 ) in 3M HCIO^ -NaClO^ solutions result in values of k c and k-^  equal to 0 . 2 5 3 M"1 hr" 1 and 0 . 0 8 9 M"1 hr" 1 respectively. That is, the rate increases as the acidity is increased. This fact had been obscured before because of medium effects. Gilks and Waind ( 6 0 ) have shown that the exchange rate decreases linearly with increase in ionic strength. Furthermore, substitution of HCIO^ for NaClO^ increases the rate at a CIO^ concentration of 1.22M, has no effect at 3 . 0 M and decreases the rate at 6 . 0 M. No convincing interpretation is available for this behaviour. It may be related to changes in activity coef-ficients or degree of hydration. The effects of chloride, bromide and cyanide ions on the exchange rate are similar. As the concentration of halide ion is increased the rate decreases, passes through a minimum and then increases sharply. This is at-tributed to complexing, the lower thallic complexes being less reactive and the higher thallic complexes and the thallous complexes being more reactive than the aquo ions. Baysal ( 6 1) has determined the rate constants in terms of reactions between various complexes. These are shown in Table I. In the case of bromide ion, exchange also occurs through the reversible oxidation of Br" by Tl(lII) ( 6 2 ) . Table I Exchange rate constants f o r t h a l l i u m chloro complexes Ionic strength 3.0 M, Temperature 30.0° Exchanging Species Rate Constant M"1 min' T l + -- T l 3 + 8.22 X IO" 3 T l + • - T 1 C 1 2 + 3.5 X 10-5 T l + • * T 1 C 1 2 + 1.06 X 10"4 T l + " TICI3 0 + T l 5.8 X T1C1 • • T l C l ^ 0.32 T1C1 2 - • ' T1C1 A" 151 22 Sulfate has a pronounced c a t a l y t i c e f f e c t on the exchange and the k i n e t i c s are not simple, Brubaker et a l (63, 6 4 ) i n t e r p r e t the r e s u l t s as i n v o l v i n g one and three (but not two) s u l f a t e groups, e.g. 1 T i - cr - s 1 0 0 « 2-0 - S" - 0 - T1(S0,) 1 4 0 The reduction of T l ( l I I ) by Fe(II) has also been the subject of several i n v e s t i g a t i o n s . The i n i t i a l rate i s f i r s t order i n each of F e ( I l ) and T I ( I I I ) , but the rate f a l l s o f f above 60% r e a c t i o n ( 6 5 ) . Ashurst and Higginson ( 6 6 ) showed that t h i s was due to i n h i b i t i o n by the product F e ( I I I ) , A two step mechanism with the formation of T l ( l l ) explains the k i n e t i c behaviour. T1(II) + Fe(III) T1(I) + Fe(III) An i n i t i a l two electron step i s rejected as the rate i s not a f f e c t e d by adding T l ( l ) . The a c i d dependence shows two paths inverse f i r s t and second order + 2-i n H , The r e a c t i o n i s accelerated by SO^ . Chloride ion i n h i b i t s the reac-t i o n , but Duke and Bornong (67) found that at higher chloride concentrations the rate again increased r i s i n g above the chloride f r e e rate and then l e v e l i n g o f f . They Interpret t h i s as being caused by complexing of the T l ( I l ) and con-clude that since the T l ( l ) - T l ( l I I ) exchange behaves i n a s i m i l a r manner i t 2+ - 0 Tl and (SDA)T1 -T I ( I I I ) + Fe(II) ^ "-1 T1(II) + Fe(II) 23. must also Involve T l ( l l ) . This does not f o l l o w as the l e v e l i n g o f f i s explained simply by assuming complete conversion of T l ( l l l ) to, say, T l C l ^ so that f u r t h e r C l " has no e f f e c t , 2+ Irvine (68) examined the oxidation of O s ( b i p ) 3 by T l ( l l l ) . I t apparently proceeds i n two steps and l i k e the F e ( l l ) r e a ction, i s i n h i b i t e d by c h l o r i d e ion. Unlike the F e ( I I ) reaction the rate increases s l i g h t l y with increased a c i d i t y , but t h i s may be a medium e f f e c t , Ashurst and Higginson (69) also studied the oxidation of T l ( l ) by C o ( I I l ) , This r e a c t i o n also appears to proceed through the formation of T l ( I l ) according to the mechanism k l Co ( I I I ) + T1(I) Co(II) + T1(II) k o C o ( l l l ) + T1(II) — C o ( l l ) + T I ( I I I ) The oxidation of H g ( l ) 2 by T l ( l I I ) studied by Armstrong and Hal-pern (70) i s best explained by a two electron t r a n s f e r mechanism. The rate law can be described by the expression - d f r l ( I I l l ) = k ' [ H g ( l ) 2 ] [ T l ( H l ) ] d t [ H g ( l l ) ] [ H + ] This i s consistent with a rate determining r e a c t i o n between T l ( I I l ) and Hg(0) formed by the H g ( I ) 2 dismutation equilibrium, 2+ s. 2+ Hg 2 Hg + Hg-" Hg + T10H2"1" Hg2"1" + T l + + OH** In contrast to other T l ( l I I ) reactions chloride i o n increases the rate, ' 2+ presumably by complexing with Hg and s h i f t i n g the equilibrium to the r i g h t . 24. There i s also a r e t a r d i n g e f f e c t of perchlorate ion which may be due to 2+ complexing with Hg 2 • Another reaction of T l ( I I I ) which may proceed by way of a two electron t r a n s f e r step i s the oxidation of C r ( I l ) . Ardon and Plane (71) found that the chromium product i s a dinuclear C r ( l I I ) species and the f o l -lowing steps are suggested C r ( I l ) + T I ( I I I ) > Cr(IV) + T l ( l ) C r ( I I ) + C r ( l V ) > C r ( I I l ) 2 A l t e r n a t i v e l y , dlmerization could precede the oxidation, e.g. 2 C r ( l l ) C r ( I I ) 2 C r ( l l ) 2 + T l ( I I I ) > C r ( I I I ) 2 + T l ( l ) A k i n e t i c study could d i s t i n g u i s h between these p o s s i b i l i t i e s . Higginson et a l (72) have examined the oxidation of V ( l I I ) and V(IV) by T l ( I I I ) . I t was thought that the V ( l l l ) r e a c t i o n might make use of a two-equivalent step. This i s a very f a s t reaction and the k i n e t i c s were not determined. However, the mechanism TI ( I I I ) + V ( I I I ) > T1(II) + V(I?) T1(II) + V(III) > T1(I) + V(IV) i s preferred over T I ( I I I ) + V(III) > T1(I) + V(V) V(V) + V(III) > 2V(IV) c h i e f l y because no V(V) i s detected. 25. The V(IV) r e a c t i o n i s slow and must be c a r r i e d out above 60'. I t apparently proceeds i n a manner analogous to the Fe(II) reaction and the f o l -lowing mechanism i s proposed T I ( I I I ) + V(IV) T1(II) + V(V) k - l T1(II) + V(I7) k 2 > T1(I) + V(V) Complications a r i s e from the slow decomposition of T l ( l l l ) , T I ( I I I ) + H 20 > T1(I) + i o 2 + 2H + The unequivocal r e s o l u t i o n of the foregoing reaction i s of some importance i n view of Sykes 1 contention (73) that i t proves the existence of a two e l e c t r o n t r a n s f e r i n the T l ( l ) - T l ( l l l ) exchange. The basis of t h i s contention i s that the T l ( l l l ) - V(IV) reaction at 80° has about the same rate as the T l ( l ) - T l ( I I l ) exchange ( l x 1 0 " % " 1 min'--'- compared to 5 x 10" 4 M""*" min"^). I t would be expected then that the addition of T l ( l ) would catalyse the i n i t i a l disappearance of TI'(III) i f the exchange rea c t i o n u t i l i z e s T l ( I l ) as an intermediate. Sykes found no change i n the rate of the V(IV) reaction when the i n i t i a l T l ( l ) content was varied from zero to 0.13 M and, therefore, concluded that the T l ( l ) - T l ( I I l ) takes place by a one step two ele c t r o n t r a n s f e r . S i m i l a r arguments have been put f o r t h by Gryder and Dorfman (74-). They reported on the oxidation of T l ( l ) by Ce(IV) i n 6 M HNO^ at 54°. The empirical rate law i s - d M l J = k o [ C e ( l V ) ] + k ^ C e d V ^ T K l ) ] The r e a c t i o n i s i n h i b i t e d by C e ( l l l ) , but not by T l ( I I I ) . 26. While Gryder and Dorfman may be cor r e c t i n assuming e i t h e r a two ele c t r o n T l ( l ) - T l ( l l l ) exchange or a two electron Ce(IV) - T l ( l ) reaction i t i s d i f f i c u l t to assess t h e i r claim because of the paucity of experimental data. There are quite l i k e l y r a d i c a l s involved i n the Ce(IV) - T l ( l ) r e a c t i o n (75) which can lead to quite complicated mechanisms. Traces of chl o r i d e would immediately destroy the argument as chlo r i d e i n h i b i t s the exchange re a c t i o n , but catalyses (40) the oxidation r e a c t i o n . Sykes 1 case i s on better grounds, but r e s t s heavily on the proposed mechanism f o r the V(IV) - T l ( l I I ) reaction. The non-metallic reductant, formic acid, appears to reduce T l ( I I l ) by a two equivalent e l e c t r o n transfer. Halvorson and Halpern (76) i n t e r -preted the k i n e t i c s as proceeding v i a a T l ( l l l ) - HCOOH complex and presented s p e c t r a l evidence to support t h i s . Subsequently Halpern and Taylor (77) showed that there i s no acid independent path so consider that the i n t e r -mediate i s a formate complex. The mechanism i s represented as T l 3 + + HCOO" HC00T1 2 + HCOOT1 2 + > H + + C0 2 + T l + F i n a l l y , Brubaker (78) has reported preliminary r e s u l t s (obtained subsequent to the publ i c a t i o n of t h i s work (79)} on the e f f e c t s of some di c a r b o x y l i c acids on the rate of the U(IV) - T l ( l l l ) r eaction. Experimental Materials Uranium solutions were prepared i n a manner s i m i l a r to that of Halpern and Smith (53). Uranyl perchlorate solutions were made by d i s s o l v i n g an uranyl s a l t , u s u a l l y the n i t r a t e or acetate, i n d i l u t e p e r c h l o r i c acid and p r e c i p i t a t i n g uranium peroxide by adding hydrogen peroxide. The peroxide p r e c i p i t a t e was then dissolved i n pe r c h l o r i c acid. The most convenient way of accomplishing the p r e c i p i t a t i o n was slowly to add 3% H 20 2 while s t i r r i n g . This gave a f i n e yellow p r e c i p i t a t e which was washed several times by decan-t a t i o n . Water was added to the p r e c i p i t a t e and the r e s u l t i n g s l u r r y was heated to b o i l i n g and allowed to cool while s t i r r i n g . This gave a p r e c i p i -tate which could be f i l t e r e d r e a d i l y using a medium sintered glass f i l t e r . The p r e c i p i t a t e was dissolved by b o i l i n g i n 3M HCIO^. The p r e c i p i t a t i o n was repeated twice to ensure the removal of fore i g n anions. I f uranium peroxide was p r e c i p i t a t e d by using 30$ H 20 2 the r e s u l t i n g p r e c i p i t a t e was very slimy and could not be f i l t e r e d . Uranium (IV) solutions were prepared by e l e c t r o l y t i c reduction of the uranyl solutions. The reduction was ca r r i e d out using a gold plated platinum gauze cathode and a platinum wire anode. The s o l u t i o n was kept cold by immersion i n an i c e bath and was s t i r r e d magnetically during the reduction. About 9056 reduction was obtained leading to stock solutions which were 0.22 M U(IV), 0.02 M U(VI) and 1.7 M HCIO^. These solutions were quite stable over a period of several weeks. When the U(IV) content dropped too low because of the i n e v i t a b l e oxidation by atmospheric oxygen, the s o l u t i o n was again r e -duced. Thallium ( I I I ) perchlorate solutions were prepared by the method of Armstrong and Halpern (70). Thallium (I) s u l f a t e was dissolved i n pe r c h l o r i c acid and was oxidized to T l ( I I l ) by sodium bromate. Excess bromine was removed by b o i l i n g and then brown t h a l l i c hydroxide was p r e c i p i t a t e d by adding sodium hydroxide. The p r e c i p i t a t e was washed several times by decantation then d i s -solved by b o i l i n g with 6056 HC10,. The p r e c i p i t a t i o n and d i s s o l u t i o n were 28. repeated twice. The resulting solution was 0.50 M Tl(IIl)s> 0.02 M Tl(l) and 8.3 M HCIO^ . Silver perchlorate was recrystallized several times from concen-trated perchloric acid. Sodium perchlorate solutions were made by neutralizing reagent grade sodium carbonate with perchloric acid. Sodium perchlorate from G. F. Smith Chemical Company gave erratic results. All other chemicals were reagent grade and were used without fur-ther purification. Water redistilled from alkaline permanganate was used in preparing a l l reaction solutions. Analytical Uranium (IV) was determined by titration with a standard solution of Ce(IV) in sulfuric acid. Osmium tetroxide was used as a catalyst and ferroin as the indicator. The Ce(lV) solution was standardized against pri-mary standard ferrous ammonium sulfate. Total uranium was determined by passing the solution through a Jones reductor putting all the uranium in the U(IV) state. The small amount of U(III) formed was quickly oxidized to U(IV) by air. The solution was then titrated with Ce(lV). Thallium (I) was titrated with a standard KBrO^ solution using methyl orange as indicator. The KBrO-j was standardized against primary standard AsgO^ .. Total thallium was determined by reduction of the T l ( l l l ) with sodium sulfite and titration with K B r 0 3 a^"'ter removal of excess SGFg by boiling. 29. P e r c h l o r i c acid concentrations were determined by t i t r a t i o n with sodium hydroxide standardized against potassium acid phthalate. Allowance was made f o r the metal content i n c a l c u l a t i n g the a c i d i t y of the metal stock solutions. K i n e t i c Measurements The majority of the reactions were conducted i n the thermostatted c e l l compartment of a Beckman DtT spectrophotometer. The k i n e t i c s were f o l -lowed by determining the l i g h t absorption at 6500 X which i s the maximum of the U(IV) ion. I t was established that Beers' law i s obeyed and that other ions d i d not i n t e r f e r e at t h i s wave length. A s o l u t i o n of the same i o n i c medium was used as the reference s o l u t i o n . In conducting an experiment the usual procedure was to prepare a s o l u t i o n of T l ( l I I ) containing appropriate amounts of HCIO^ and NaClO^ and to bring t h i s s o l u t i o n to the desired temperature i n a thermostatted water bath. A known volume of the U(lV) s o l u t i o n , u s u a l l y one ml., was then added to a known volume of the T l ( I I I ) s o l u t i o n u s u a l l y about 50 ml. The r e s u l t i n g reaction s o l u t i o n was quickly mixed and a portion transferred to a one cm. absorption c e l l . The o p t i c a l density was then recorded at predeter-mined i n t e r v a l s . Zero time was taken as the time of mixing. The operations of mixing, r i n s i n g and f i l l i n g the c e l l and placing i t i n the spectrophoto-meter took about 15 seconds. The absorbency at zero time was determined by running a blank with a l l components except the th a l l i u m s o l u t i o n . The temperature i n the c e l l compartment was c o n t r o l l e d to ±0.1° C. Some experiments were conducted by allowing the reaction to take place i n f l a s k s contained i n a thermostatted water bath. Samples were with-drawn from time to time and the o p t i c a l density measured at 6500 X. I d e n t i c a l r e s u l t s were obtained with both procedures. 30. Results and Discussion Assuming that the rea c t i o n has the stoichiometry represented by U ( I V ) + T I ( I I I ) > U(V1) + T1(I) and assuming also that the rea c t i o n i s f i r s t order i n both U(IV) and T l ( l l l ) then the rate law takes the usual form (80), H = k(er - x) (b - x) This integrates to l o g sfLzuc) = k t b - a? a b(a : — x) Here x i s the decrease i n reactant concentration and a and b are i n i t i a l concentrations. In the present case where the reaction i s followed by measuring the o p t i c a l density of the s o l u t i o n and U(IV) i s the only absorbing species the integrated expression may be written as b - a b/a D or Jt=~b l D g a/b (D - Doo) In these expressions a and b are the i n i t i a l concentrations of U(IV)) and 11(111)'; r e s p e c t i v e l y . D i s the o p t i c a l density at time t , D 0 i s that at zero time and 1^ i s that f o r the completed reacti o n . The f i r s t expression corres-ponds to T l ( I I l ) i n excess and the second to an excess of TI(IV). The f i r s t expression i s the most convenient to use and, consequently, most experiments 31. were done with the i n i t i a l T l ( l I I ) concentration 2 to 3 times that of the i n i t i a l TJ(IV) concentration. The stoichiometry was checked by two procedures. In several experi-ments the reaction was allowed to go to completion. In those cases i n which T l ( l l l ) was i n excess, the f i n a l T l ( l ) concentration was determined by t i t r a t i o n with KBrO^. A f t e r allowing f o r the T l ( l ) present i n i t i a l l y , i t s f i n a l concentration was found to be e s s e n t i a l l y equal to the i n i t i a l concen-t r a t i o n of tf(lV). When U(lV) was i n i t i a l l y i n excess, the amount consumed was found to be the same as the i n i t i a l concentration of T l ( l I I ) , These r e s u l t s are shown i n Table II., The second order nature of the reaction i s demonstrated i n Figure 2. Both the o p t i c a l density and the second order integrated function have been plotted. The f a c t that the r e a c t i o n went e s s e n t i a l l y to completion with no appreciable deviation from the l i n e a r second order p l o t suggests that the products U(VT) and T l ( l ) have no e f f e c t on the r e a c t i o n . Further confirmation of the nature of the r e a c t i o n i s given by the data of Table I I I . Here i t i s shown that v a r i a t i o n of U(IV) from 2.4 x 10~ 3 M to 16.7 x-:10"3 M and of T l ( I I I ) from 5.0 x 10" 3 to 19.7 x I O - 3 M had no e f f e c t on the second order rate constant. In view of the e f f e c t of C l " on the reactions of T l ( I I I ) i t s influence was i n v e s t i g a t e d here. The reagent was NaCl. The r e s u l t s are shown i n Table IV and Figure 3. As the c h l o r i d e concentration was increased the rate decreased and reached a minimum at a C l " / T l ( l l l ) r a t i o of 2. Further additions of NaCl increased the rate very s l i g h t l y . I t seems very l i k e l y that the reduction i n rate i s due to the formation of T l ( I I l ) complexes, e.g. T1C1 and T1C1„ . 32. Table II Stoichiometry of U(IV) - Tl(IIl) reaction HlV)± T K I I I ^ U(IV) f Tl(I) f (a) M x 103 M x 103 M x IO3 M x IO3 R (b) 5.54 15.0 - 5.73 1.04 4.40 15.0 - 4.57 1.04 4.77 16.7 - 5.00 1.05 4.77 16.7 - 5.04 1.06 3.55 20.6 - 3.55 1.00 5.07 15.5 - 5.22 1.03 4.10 10.1 - 4.17 - 1.02 4.10 10.1 - 4.12 1.00 4.10 10.1 - 4.07 0.97 4.10 10.1 - 4.14 1.01 4.10 10.1 - 4.15 1.01 4.08 9.0 - 4.09 1.00 3.77 9.4 - 3.79 1.01 3.77 9.4 - 3.75 0.99 Ave. 1.02 9.44 6.92 2.91 - 1.06 16.70 6.92 10.07 - 1.04 9.61 5.01 4.71 - 1.02 9.96 6.02 4.05 - 1.02 10.83 5.02 5.80 - 1.00 9.87 5.48 4.41 - 1.00 Ave. 1.02 (a) corrected for i n i t i a l Tl(l) (b) for excess T l ( l l l ) R = T1(I)f/JJ(lV)1 for excess U(IV) R = T l l I I I ^ / f l K l V ^ - TJ(IV)f) Table I I I E f f e c t of i n i t i a l concentrations of U(IV) and T l ( I I l ) on second order rate constant of U(IV) - T l ( I I I ) r e a c t i o n Temperature 2 5 . 0 ° U(IV) T I ( I I I ) k M M x l O 3 M x 1 0 3 M ' ^ i n " 1 3 7 . 6 17 . 0 0 . 5 8 3.9 17 . 0 0 . 5 3 9 . 4 6 . 9 0.72 16.7 6 . 9 0 . 6 7 3 . 9 1 0 . 0 0 . 5 9 3 . 6 1 0 . 0 0 . 5 9 (a) 1 0 . 0 5 . 0 0 . 5 8 9 . 6 5 . 0 0 . 5 9 (a) 1 0 . 0 6 . 0 ' 0 . 6 1 2 . 8 8 5.5 1 5 . 0 0 . 5 6 4 . 4 15 . 0 0 . 5 4 1 0 . 8 5 . 0 0.55 9.9 5.5 0 . 6 0 2.82 4 . 3 7.7 0 . 6 0 4.1 9 . 0 0.63 (b) 3 . 8 9 . 4 0 . 5 9 (b) 2 . 4 9 . 0 0 . 5 6 (b) 2 . 9 4 4 . 0 19.7 0 . 5 2 (b) Contained 3 . 3 x 1 0 ' % U(VI) Reactions i n spectrophotometer, a l l others i n f l a s k s and samples withdrawn f o r analyses 35. Table IV Effect of NaCl on rate of II(IV) - Tl(III) reaction Temperature 25.0° HCIO^ 2.83 M &(IV) M x 10" TI(III) M xlO" NaCl M x 10" M'^ -min"1 4.1 4.1 4.1 4.1 4.1 4.1 3.2 4.0 4.0 4.0 4.0 4.0 4.0 9.0 9.0 9.0 9.0 9.0 9.0 9.0 9.5 9.5 9.5 9.5 9.5 9.5 0.0 1.5 2.9 5.9 8.8 11.8 H.1 14.7 17.6 20.6 23.5 29.4 32.3 0.629 0.553 0.414 0.297 0.124 0.054-0.039 0.034 0.021 0.04.7 0.061 0.066 0.065 36. 37. The f a c t that there i s no pronounced increase of rate beyond the minimum ind i c a t e s that the higher chloro complexes are also unreactive. This con-t r a s t s with the F e ( l l ) and T l ( l ) reactions. The behaviour of s u l f a t e , added as NagSO^, was also studied. As noted i n Table V and Figure 4 s u l f a t e exerts a c a t a l y t i c e f f e c t . At the higher s u l f a t e concentrations there i s considerable s c a t t e r i n the r e s u l t s . However, the general trend i s i n accord with an a d d i t i o n a l path which i s f i r s t order i n s u l f a t e ion. Table VI l i s t s values of the second order rate constant obtained 2+ 2+ when metal perchlorates were added. Neither Cu nor Hg had any e f f e c t . Small amounts of A g + reduced the rate by about 30$, but no f u r t h e r reduction i n rate was accomplished by i n c r e a s i n g the amount of s i l v e r perchlorate. The reason f o r t h i s curious behaviour i s not known. These cations were chosen because of t h e i r e f f e c t on the rate of oxidation of U(IV) by 0 2 (53). Both C u 2 + and H g 2 + catalyse t h i s r e a c t i o n , whereas Ag + i s an i n h i b i t o r . The e f f e c t of v a r i a t i o n of the i o n i c strength on the rate of the r e a c t i o n was examined. This was done by holding the p e r c h l o r i c a c i d con-centration constant at 1.76 M and adding varying amounts of sodium perchlorate. The data f o r these experiments are contained i n Table VII. The second order rate constant increases as the concentration of sodium perchlorate i s increased as i s shown i n Figure 5. There i s no evidence f o r complexing of e i t h e r U(IV) or T l ( l l l ) with C10^" and a medium e f f e c t i s the most l i k e l y explanation. The e f f e c t i s i n the d i r e c t i o n predicted by the Debye-Huckel theory f o r a r e a c t i o n between two ions of l i k e charge. However, medium e f f e c t s i n solutions of high i o n i c strength are not q u a n t i t a t i v e l y predic-t a b l e . As shown i n Figure 5, l o g k v a r i e s l i n e a r l y with the concentration of NaClO^, I t i s not c l e a r whether or not t h i s has any s i g n i f i c a n c e . Table V Effect of NagSO^  on rate of U(IV) - Tl(III) reaction Temperature 25.0° TJ(IV) 3.19 x 10*% HCIO^ 2.83M H.(III) 9.04 x 10"% N a ^ k M x 103 M"1min"1 0.0 0.59 1.7 1.10 3.5 1.59 5.2 2.O4 7.0 2^30 7.8 3.06 8.7 2.95 9.6 3.16 10.5 2.80 11.3 3.37 14.0 3.61 39. 40. Table VI Effect of some metal ions on rate of U(IV) - T l ( l l l ) reaction Temperature 2 5 . 0 ° U(IV) 3.5 x 10-3M HCIO^ 2.82 M Tl(III) 9.0 x 10"3M Concentration ^ Added Salt M x 10 4 M"1 min - 1 0.0 0.59 Cu(C10^)2 20 0.57 Hg(C10^)2 12 0.61 A&CIO^ 4.2 0.54 8.4 0.50 16 0.42 16.8 0.44 21.0 0. 48 21.0 0.42 33. 7 0.42i 34.0 0.38 52.6 0.42 68..1 0.38 84.1 0.41 Table VII Effect of NaC104 on rate of U(IV) - Tl(IIl) reaction Temperature 25.0° U(IV) 3.0 x 10-3M HCIO^ 1.76 M TI(III) 9.0 x 10" 3M NaCIO 0 O.84 1.29 1.71 2; 58 2.58 2.96 M"1min-: O.64 0.92 1.20 1.34 1.85 1.94 2.30 42. 43. The rate of the reaction was found to depend strongly on the concentration of HCIO^. Tables VIII to XII show the e f f e c t of the v a r i a t i o n of HC10. on the second order rate constant at 16°, 20° and 25°. At the low 4 acid concentrations the rate becomes quite f a s t and s e n s i t i v e to impurities. There was considerable scatter of the low acid r e s u l t s and t h i s was traced, i n part, to the NaClO^ used i n adjusting the i o n i c strength to a constant value. The v a r i a t i o n of rate i s i l l u s t r a t e d i n Figure 6 which shows rate plots obtained at d i f f e r e n t HCIO^ concentrations and constant i o n i c strength. These plots are t y p i c a l i n that they show a small, p o s i t i v e , ordinate i n t e r c e p t . This i n d i c a t e s that there may have been a systematic er r o r i n the procedure. No attempt was made to remove dissolved oxygen and t h i s or other impurities may have contributed. Local high concentrations during the i n i t i a l mixing of reactants would lead to such zero time errors. The f a c t that the rate increases as the a c i d i t y i s decreased suggests a rate law of the form - d M l V } l = k [ U ( l V ) ] [ T K H l ) ] f H + j " n Accordingly the second order rate constant, k, was p l o t t e d against [ H + J on a logarithmic p l o t . This i s shown i n Figure 7 and y i e l d e d a s t r a i g h t l i n e with a slope of -1.39. Since t h i s value of n l i e s between 1 and 2 the rate law was written At the higher a c i d i t i e s both U(IV) and T l ( I I I ) e x i s t l a r g e l y as the simple aquo ions U 4 + and T l 3 + . As the a c i d i t y i s lowered some hydrolysis Table VIII Effect of HCIO^  on rate of U(IV) - Tl(lII) reaction Temperature 25.0° Ionic strength maintained at 2.88M with NaClO^ prepared from NaOH and HCIO^ TI(III) U(IV) HCIO^ k MxlO3 MxlO3 M M"1 min"1 10.1 4.1 0.478 5.93 10.1 4.1 0.752 3.16 10.1 4.1 1.026 2.14. 10.1 4.1 1.300 1.54 10.1 4.1 1.57 1.21 10.1 4.1 1.85 1.01 20.1 4.1 0.698 2.83 20.1 4.I 0.754 2.98 20.1 4.1 0.808 2.70 20.1 4.1 2.29 0.70 20.1 4.1 2.56 0.62 20; 1 4.1 2.94 0..52 Table IX Effect of HCIO^ on rate of TJ(lV) - Tl(III) reaction Temperature 25.0° Ionic strength maintained at 2.82M with NaCIO 4 prepared from NapCXD- and HC10. TI(III) U(IV) HCIO^ k M x l O 3 M x 103 M M'-'-min"1 9.0 2.7 0.477 6.70 9.0 2.7 0.565 4.66 9.0 2.7 0.740 3.00 9.0 2.7 0.913 2; 15 9.0 2.7 1.44 1.23 NaClO^ from NaHC03 and HCIO^ 9.0 2.4 0.567 4.86 9.0 2.4 0.744 3.20 9.0 2 .4 0.917 2; 53 Table X Effect of HCIO^ on rate of U(IV) - T l ( l l l ) reaction Temperature 25.0° Ionic strength maintained at 2.82M with NaClO^ prepared from recrystallized NagCO^  and HCIO^ TI(III) U(;IV) HCIO^ k M x l O 3 M x l O 3 M M'^mln"1 9.04. 2; a 0.4-67 7.02 9.04. 2.41 0.556 5.45 9.04 2.41 0.731 3.74 9.04 2.41 0.906 2.78 9.04 2.41 1.43 1.40 9.04 2;41 2.82 0.563 Table XI Effect of HCIO^ on rate of TJ(IV) - Tl(III) reaction Temperature 20.0° Ionic strength maintained at 2,82 M with NaClO^ prepared from recrystallized NagCO^  and HCIO^ TI(III) U(iy) HC10. k 4 M x 103 M x 103 M M^min"1 9.04. 2,20 1.43 0.703 9.04 2.20 0.906 1.54 9.04 2.20 0.556 2.97 9.04 4.12 1.08 1.12 9.04 4.12 0.731 1.98 9.04 4*12 0.467 3.87 9.04. 3.42 2.83 0.294 Table XII Effect of HCIO^ o n r a t e o f T T^ I V^ " T 1 ( I I : [ ) Temperature 16.0° Ionic strength maintained at 2.82 M with prepared from recrystallized NagCO^  and TI(III) U(IV)) HCIO^ k M x 103 M X I O L ? M M~1min" 9-04 3.88 0.487 2.63 9.04 3.88 0.556 1.73 9*04 3.88 0.731 1.15 9.04 3.88 02906 0.878 9.O4 3.88 1.08 0.634 9.04 3.88 1.43 0.425 9.04 3.88 2.82 0.158 reaction NaClO^ HC10, Figure 6. Second order rate plots at 25.0 for various HC10, concentrations. Ionic strength adjusted to 2.8M with NaClO^ a = [U(IV)] Q = 2 . 4 x 10"3M; b = [ T l ( I I I ) J o = 9.0 x 10"3M Time min. 50. 51. w i l l occur' r e s u l t i n g i n the production of U0H 3 + and TIOH2* ions. I f , as i s l i k e l y , these are the only species of importance one can write for the t o t a l concentrations fr(iv)] = = [u 4 +] + [UOH 3 +] and [TI(III)] = = [ T I 3 + ] + [TIOH 2 +] Then L V + ] • -1 and [ T l 3 + J = = [TI(III)] / f l + K T 1 [ H + j -1 4+ 3+ where and K are the hydrolysis constants of u and Tl . The rate law then becomes -dfo(IV)] = [U(IV)][T1(III)] ( k j H + J + k,) dt ' -±-r ~ ([H +J +K t J )X[H +] + K n ) In terms of the experimental second order rate constant k = k j H * ] + + ^ m * ] + v In order to determine the constants k^ and kg the l a s t equation was written k ([H +] + y d X 1 " ] + K T 1) = k^H"*] + kg and the l e f t hand side was plotted against ^ H +J as shown i n Figure 8. The slope and intercept of t h i s plot y i e l d k^ and kg respectively. The values of and which were used are given i n Table X I I I , The values at 25° were taken from the l i t e r a t u r e (81, 82). The values at the other temperatures were calculated using an enthalpy of hydro-l y s i s of 11.0 kcal/mole. This corresponds to the measured value f o r 52. (83, 84) and i s probably a reasonable estimate f o r T l as most known heats of hydrolysis of metal ions l i e i n the range 9 - 1 2 kcal/mole (85). Fortun-a t e l y the extent of hydrolysis of both ions i s small so that the k i n e t i c r e s u l t s are not s e n s i t i v e to errors i n the hydrolysis constants. The rate constants k-^  and kg are l i s t e d i n Table XIV together with the a c t i v a t i o n parameters. These parameters were ca l c u l a t e d from the Arrhenius plots shown i n Figure 9. The comparatively narrow temperature range was imposed by experi-mental conditions. The lowest temperature was that which could be conven-i e n t l y maintained i n the spectrophotometer. At temperatures higher than 2 5 ° a large v a r i a t i o n i n a c i d i t y could not be obtained because the rates i n low acid solutions were too f a s t to follow. The reaction was followed up to 40° at a HCIO^ concentration of 2,83 M. The observed second order rate constants, k, are contained i n Table XV together with values c a l c u l a t e d from k^ and kg. I t i s seen that the agreement i s very good. The k i n e t i c evidence i s consistent with a simple one stage, two electron t r a n s f e r step, TJ(IV) + T I ( I I I ) U(VI) + T1(I) The case where U(V) and T l ( I I ) are formed i n the rate determining step and then react with each other before d i f f u s i n g from the "solvent cage" must be considered as k i n e t i c a l l y equivalent to the one step mechanism. The a l t e r n a t i v e i s a sing l e electron t r a n s f e r step which must then be followed by various f a s t reactions, e.g. U(IV) + T I ( I I I ) — ^ U(V) + T1(II) 53. 1.0 2.0 H Figure 8. E f f e c t of a c i d i t y on rate at various temperatures, Ionic strength 2.9 M Table XIII Z+ 3+ Hydrolysis constants of U and Tl Temp % °n M" M T 1 2 5 0.021 0.073 20 0.016 0.056 16 0.012 0.041 54 Table XIV Kinetic data for the oxidation of U(IV) by Tl(III) Ionic strength 2.82 M o o ^ F ± ^ H * ^ S ± 16.0 20.0 25.0° kcal/mole kcal/mole e.u. k-LxlO^sec'1 0.57 1.02 2. XI 19.7 24.6 16.4 2 -1 kgxlO ,Msec 0.67 1.17 2.13 19.7 21.7 6.7 55. o r. J . J 10 /T A Figure 9. Arrhenius plots f or rate constants ki and k2 of the U(IV) - T l ( I I I ) reaction Table XV Variation of rate of U(IV) - T l ( l l l ) reaction with temperature HC104 2.83 M tJ(iv) TI(III) Temp k k calculi ! x 103 M x 103 °C M"1 min""1 M"1 min" 3.42 9.04 15.0 O . I 4 O 3.42 9.04 15.0 O . I 4 O 3.42 9.04 20.0 0.294 -3.42 9.04 20.0 0.294 3.77 9.40 25.0 0.588 -3.77 9.40 25.0 0.588 3.27 9.05 30.0 1.16 1.12 3.27 9.05 30.0 1.16 3.28 9.05 35.0 2.16 2.11 3.28 9.05 35.0 2.15 3.28 9.05 40.0 3.84 3.89 3.28 9.05 40.0 3.62 (a) Calculated using k x and kg from Table XIV and and K, from Table XIII. 57. U(V) + T I ( I I I ) > TJ(VI) + T1(II) U(IV) + T1(II) > U(V) + T I ( I I I ) 2U(V) > U(IV) + U(VI) Such a scheme seems l e s s p l a u s i b l e f o r these reasons? 1) The k i n e t i c s were found to be quite straightforward with no 2+ suggestion of competing steps. The insensitiveness of the rate to Cu and 2+ Hg suggests that no active intermediates are present. The anomalous behaviour of A g + i s neither pronounced enough nor of the type to be expected from a chain mechanism. 2) The formation of two unstable species as the r a t e determining step would probably require a high a c t i v a t i o n energy. The enthalpy of f o r -mation of T l ( I l ) i s not known, but an upper l i m i t can be set from the data of F e ( l l ) - T I ( I I I ) r e a c t i o n (65). The rate determining step i s F e 2 + + T l 3 + » F e 3 + + T l 2 + The enthalpy of r e a c t i o n i s not known, but can be no higher than A H * which i s 17 kcal/mole. Using the known values (86) f o r the other e n t i t i e s t h i s leads to an upt>er l i m i t of A H ^ f o r T l 2 + of 55 kcal/mole. I t i s probably not much l e s s than t h i s , the corresponding value f o r T l i s 4? kcal/mole. The enthalpy of r e a c t i o n f o r U 4 + + T l 3 + + H 20 > TJ0 2 + + T l 2 + + 4H + can now be determined as 44 kcal/mole, again an upper l i m i t . The corres-ponding value f o r the complete re a c t i o n U 4 + + T l 3 + + 21^0 > U 0 2 2 + + T l + + 4H + i s -12.5 kcal/mole. Thus, on energetic grounds the one step r e a c t i o n i s 58. favoured. This i s also i n l i n e with the observations of Newton and Rabideau (87) that f o r actinide elements, lower enthalpies of reaction correspond to lower enthalpies of a c t i v a t i o n . This argument m i l i t a t e s against the f o r -mation of U(V) and T l ( l l ) as separate e n t i t i e s and thus supports a true single step mechanism. Single step r e f e r s only to the electron t r a n s f e r or valence change process. There may of course be other reactions necessary f o r the products to assume t h e i r c o r r e c t configurations. 3) In both the F e ( l l ) - T l ( l I I ) and the C o ( I I l ) - T l ( l ) reactions which produce T l ( l l ) as an intermediate, the subsequent r e a c t i o n of T l ( l l ) i s slow enough to a f f e c t the k i n e t i c s because the reverse of the i n i t i a l r e a c t i o n then becomes important. No such behaviour i s observed i n t h i s case. I t i s c l e a r that the one step process does correspond with the observed behaviour. I t cannot be said to be proven and i t i s doubtful that such proof i s possible. On the other hand i t could be disproven by the observance of T l ( l l ) i n the r e a c t i o n . As mentioned before, there i s no way of d i s t i n g u i s h i n g between a sin g l e stage, two e l e c t r o n t r a n s f e r and successive one electron transfers which take place i n the "solvent cage". There i s , however, a fundamental d i f f e r e n c e i n that the l a t t e r s i t u a t i o n corresponds to two separate events. I t i s i n t h i s context that claims f o r the proof of a two electron t r a n s f e r must be judged. Thus the contentions of Gryder and Dorfman (7U) and of Sykes (73) that the T l ( l ) - T l ( l l l ) exchange reaction does not involve T l ( l l ) r e a l l y mean that a free T l ( I I ) i s not formed. The evidence concerning the composition of the activated complexes and whether or not inner or outer sphere mechanisms are involved can now be examined. 59. There are two major (and perhaps only) activated complexes i n per-chlorate s o l u t i o n s . These are k U 4 + + T l 3 + + H,0 — ( U . T l ' O H ) + H + and T J 4 + + T l 3 + + B^O (U'T1'0) 5 + + 2H + This manner of de p i c t i n g activated complexes was introduced by Newton and Rabideau (87) and has much to commend i t . I t states simply that one complex i s formed from Tr* and T l with the l o s s of a proton and the other by l o s s of two protons. There must be the residue of at l e a s t one water molecule. There may, of course, be other water molecules present. The protons may have been l o s t as the com-plexes were formed or by p r i o r hydrolysis. The k i n e t i c s t e l l s only the com-po s i t i o n of the activated complexes with respect to components whose concen-t r a t i o n s can be varied. Not much can be said with respect to the a c t i v a t i o n parameters. The parameters of both paths are very s i m i l a r with path 1 having a somewhat higher , but a more p o s i t i v e A S than path 2. The u n c e r t a i n t i e s are such that these d i f f e r e n c e s may not be r e a l . Newton and Rabideau postulated that the charge of the activated complex was the most important f a c t o r i n determining i t s entropy. That i s , a complex with a charge of -+6 i s about 20 e.u. more negative than one with a charge of +5. This i s the opposite of the present case. Complex 1 with a charge of has an entropy of -87 e.u./ mole while complex 2 has a charge of +5 and an entropy of -96 e.u./mole. I f path 2 i s written as n 4 + + n 3 + + 2 H 2 O > ( U.T1.0H.0H) 5 + + 2H + 60. the entropy of the complex becomes -80 e.u./mole. This shows why interest is taken i n Z i S ± values. The complexes (TJ.T1.0)5+ and (TJ.T1»0H«0H)5+ are different (when they imply, say, an 0 group in one and two OH groups in the other) but are not experimentally distinguishable. The entropy of acti-vation is experimentally accessible and thus might allow a choice between two activated complexes, but not between alternative formulations of the same complex. The difficulty here is that there is no standard to which values can be compared. There can be no real choice between inner and outer sphere Z.+ 2+ mechanisms. Since the IT*" ion acquires two oxygen atoms in becoming UOg i t might be argued that a bridged complex would more easily allow this. This, 'however, is not necessary as the loss of two electrons from the uranium atom would cause a strengthening of the ligand bonds and a loss of protons, e.g., HgO - U 4 + - OHg - T 2 ^ o = U = 02**" + 4H+ The catalytic effect of sulfate may be due to a bridged structure 0 0 / \ / \ 5+ U S- XT1 \ / \ / 0 0 and subsequent group transfer. On the other hand, i t s effect may be to make orbitals available which facilitate electron transfer. This may or may not require bridging. The catalysis of the reaction by 0H~ is most easily interpreted as being due to the H + released on going from U 4 + to flOg2"*", This must neces-sarily occur whether or not a bridged mechanism is involved. The Tl(l) -61. T l ( I I l ) exchange which does not involve a change of structure i s only s l i g h t l y affected by OH". The i n h i b i t i o n of the rea c t i o n by chl o r i d e i s at t r i b u t e d to com-plexing with T l ( l l l ) , There appear to be two f a c t o r s of importance which are again unrelated d i r e c t l y to inner or outer sphere mechanisms. The great s t a b i l i t y of T l ( l l l ) chloro complexes makes T l ( I I l ) much more r e s i s t a n t to reduction i n a thermodynamic sense. The other f a c t o r i s the l a r g e r reorgani-zation or Franck-Condon b a r r i e r which a r i s e s from the r e l a t i v e s t a b i l i t i e s of the T I ( I I I ) and T l ( I ) complexes. These f a c t o r s also a f f e c t the T l ( l ) -T l ( l I I ) exchange and are markedly reduced when the chlo r i d e concentration i s s u f f i c i e n t to cause appreciable complexing of T l ( l ) . In the U(IV) - T l ( l I I ) r e a c t i o n , increasing the chloride concentration beyond f u l l complexing of the T l ( I I I ) does not increase the rate because of the s t r u c t u r a l d i f f e r e n c e between U(lV) and U(VI). An attempt was made to reduce H g ( l l ) with U(IV), but t h i s r e a c t i o n i s too slow to follow. At f i r s t s i g h t t h i s i s puzzling since the r e a c t i o n U 4 + + 2 H g 2 + + 2 ^ 0 > U 0 2 2 + + H g 2 2 + + 4H + has a favourable free energy change of -27 k c a l . However, the rea c t i o n U 4 + + H g 2 + + 2 ^ 0 > U 0 2 2 + + H g a q + 4H + has a ^  F° of -16 k c a l which while favourable i s much l e s s than the -42 kc a l f o r the analogous T J 4 + - T l 3 + r e a c t i o n . 62. PART II KINETICS OF THE HOMOGENEOUS OXIDATION OF CARBON MONOXIDE BY METAL IONS Introduction While carbon monoxide takes part i n many chemical reactions i t i s generally considered to be i n e r t toward common o x i d i z i n g agents. The recent successes i n demonstrating the a b i l i t y of c e r t a i n metal ions and complexes to react homogeneously with hydrogen i n s o l u t i o n (88, 89) suggested that p a r a l l e l studies with carbon monoxide might prove f r u i t f u l . The oxidation of carbon monoxide d i f f e r s i n an important aspect from that of hydrogen. Carbon mon-oxide i s oxidized i n the c l a s s i c a l sense by acquiring an a d d i t i o n a l oxygen atom while the oxidation of hydrogen involves the s p l i t t i n g of the molecule with l o s s of electrons. Carbon monoxide i s the simplest, stable, heteronuclear, diatomic molecule. I t i s one of the few compounds of carbon e x h i b i t i n g an apparent valence of two. For these reasons the structure and properties of carbon monoxide have received much attention. The physical properties of carbon monoxide are remarkably s i m i l a r to those of nitrogen (90). Langmuir (91) has at t r i b u t e d t h i s to the f a c t that these molecules are i s o s t e r i c , having the same number of atoms and electrons. Mullikan (92) has assigned both molecules the same configuration KK(zcr) 2(yo-) 2(xo-) 2(w7r) 4. M o f f i t t (93) and Sahni (94) have given t h e o r e t i c a l treatments of carbon monoxide which are i n accord with a t r i p l y bonded structure with a small dipole moment d i r e c t e d toward the carbon atom. Recently, J a f f e and Orchin (95) have given a more q u a l i t a t i v e 63. discussion emphasizing hybridization effects. This latter treatment is more satisfying in that i t provides an explanation of the chemical differences between carbon monoxide and nitrogen. Much attention has been directed toward the structure and properties of the transition metal carbonyls. Mond (96) discovered these compounds in 1890, but only in the last few years has appreciable progress been made in establishing their structure. Richardson (97) and Chatt, Pauson and Venanzi (98) have summarized the current status. In the pure carbonyls the metal atom attains the effective atomic number of the next inert gas. This is borne out by the absence of monomeric carbonyls of metals with odd atomic number. The implication of the rule is that outer d orbitals are not used in bonding. This finds confirmation in the tetrahedral structure of Ni(CO) 4 3 which must use sp hybrid orbitals. The lone pair on the carbon atom is used in forming sigma bonds, but these are relatively weak. Stabilization occurs, in part, because the empty pi orbitals of carbon monoxide can accept electrons from the metal thereby preventing the accumulation df negative charge on the metal. Dicobalt octacarbonyl has achieved considerable importance in the hydroformylation of "Oxo"1 process. The process has been reviewed by Wender et al (99, 100). The mechanism may be described by the sequence Co 2(C0) 8 + Hg « > 2HCo(C0)^ HCo(CO)^ + RCH=CH2 > RCHgCHgMCO)^ RCH2CH2Co(C0)^ + CO > RCH2CH2C0Co(C0)'^ RCH CH COCo(CO) + HC'o(CO), > RCHgCHgCHO + Co 2(C0) 8 64. The step leading to the production of the aldehyde might also be RCHgCHgCOCoCCO)^ + ^ > RCh^CHgCHO + HCo(CO)^ Carbon monoxide exerts an i n h i b i t i n g e f f e c t and t h i s may be due to the reaction Cb 2(C0)g + CO > C o 2 ( C 0 ) 9 which reduces the concentration of HCo(CO)^. The feature which i s of i n t e r e s t here i s the i n s e r t i o n of CO between a metal-carbon bond. There i s some doubt as to whether t h i s i n s e r t i o n pro-ceeds d i r e c t l y or through one of the CO groups already attached to the cobalt. In the r e v e r s i b l e carbonylation of methyl manganese pentacarbonyl, CH,Mn(CO)c + CO * CR"oCOMn(CO )„ J p * J 5 the a c e t y l carbonyl group i s not that which enters from the gas phase. This was shown by C o f f i e l d et a l (101) by using C ^ tracer. Calderazzo and Cotton (102) believe that the re a c t i o n proceeds by d i r e c t combination of CO and CH 3Mn(C0)^, but have not presented a d e t a i l e d mechanism. The heterogeneous oxidation of carbon monoxide has been the subject of many i n v e s t i g a t i o n s . The topic has been reviewed by Katz (103) and Dixon and Longfield (IO4). Only i n the high temperature reduction of metal oxides i s carbon monoxide oxidized without the use of oxygen. In the majority of the low temperature c a t a l y t i c processes studied, metal oxides have been used as c a t a l y s t s . The course of the oxidation can proceed e i t h e r through r e a c t i o n of oxygen and carbon monoxide while one or both reactants are chemisorbed or through reduction of the metal oxide by carbon monoxide followed by reoxidation of the c a t a l y s t by oxygen. Both types of oxidation have been observed with the second type being favoured at higher temperatures. The low temperature eata-65. lysts are generally p-type or electron deficient semi-conductors. The homogeneous, solution oxidations of carbon monoxide have received very l i t t l e attention. There was a flurry of activity around the turn of the century perhaps prompted by the discovery of the metal carbonyls or by attempts to utilize the favourable energy change involved. In any event l i t t l e work of a quantitative nature was done. Among the metals which are reportedly produced by homogeneous reduction of aqueous solutions of their salts by carbon monoxide are gold, rhodium, ruthenium, iridium, platinum and palladium (105, 106, 107, 108, 109). The rate of the reaction with palladium is fast enough so that i t was adopted as an analytical method for carbon monoxide (110, 111). The production of silver from ammoniacal solutions of silver nitrate has also been noted (107, 112). In most of these cases the reactions may have been complicated by heterogeneous reduction, Phillips (108) reported what appears to be a homo-geneous reduction of platinum (IV) chloride to a lower platinum chlorides metallic platinum was deposited only after a considerable period of time. Several investigators have reported on the oxidation of carbon mon-oxide by potassium permanganate (105, 109, 113 )i Of these, Just and Kauko (113) have presented the only reasonably detailed kinetic study of any of the homogeneous oxidations of carbon monoxide, Hofmann ( I I 4 ) patented the oxidation of carbon monoxide by chromic acid to which mercuric oxide had been added. Mermet (115) reported that a solution of potassium permanganate acidified with nitric acid and containing silver nitrate was decolourized by carbon monoxide. Just and Kauko (113) determined the rate of reduction of perman-ganate by carbon monoxide in neutral solution at 15° and 25°. They found that 66. the reaction was f i r s t order i n both permanganate and carbon monoxide. Their r e s u l t s give rate constants of 0.070 M - 1 s e c " 1 at 15° and 0.172 M 1 sec 1 when the rate law i s expressed as The a c t i v a t i o n energy i s 14.7 kcal/mole. Bauch, Pawlek and P l i e t h (116) have reported the r e s u l t s of a k i n e t i c study of the reduction of Ag(l) and C u ( l l ) by carbon monoxide i n su l f a t e solutions. The Ag(l) reaction was c a r r i e d out at temperatures of 70 - 110° and CO pressures of 5 to 50 atmospheres. The rate was found to be second order i n Ag(I) and f i r s t order i n CO. The r e s u l t s were f i t t e d to the rate law - D ^ ( T ) J = 1.3 x 1 0 6 [ A g(I ) 3 2P C 0 exp(-l4,000/RT> M min" 1 When ammonium acetate was added the rate increased and the rate law was " d f f i ^ 1 ^ = 6.0 x 1 0 4 [ A g ( l ) J 2 P C 0 exp(-9300/RT)i M min" 1 The rate increase was a t t r i b u t e d to the increase i n pH, but the mechanism f o r both conditions was supposed to be the same, A g + + CO AgC0"+ AgC0 + + A g + A g 2 C 0 2 + A g 2 C 0 2 + + H 20 —k-> 2A"g + C0 2 + 2H + 2+ The rate determining step being the hydrolysis of the Ag 2C0 complex. I t i s not c l e a r i f t h i s i s meant to be r e a l or schematic. 67. The reduction of Cu(.II) was much slower and was done at 160 - 190° and 10 - 4.0 atm. of CO i n s u l f a t e solutions. The rate expression was - d [ C u ( l l ) ] = 2.6 x 1 0 1 3 [ C u ( l l ) ] P c 0 exp(-33,500/RT) M min" 1 The work described i n t h i s t h e s i s i s concerned with the k i n e t i c s of the oxidation of carbon monoxide by Hg(Il) and permanganate. Preliminary t e s t s were conducted by bubbling carbon monoxide through aqueous p e r c h l o r i c acid solutions of Cu(II), A g ( l ) , H g ( l l ) , F e ( I I l ) , T l ( I I l ) , C r 2 0 y 2 " and MnO/. Only H g ( l l ) and MnO^ were reduced and, consequently, d e t a i l e d studies were confined to these ions. In no cases were temperatures over 80° or pressures greater than one atmosphere used. Research undertaken i n t h i s laboratory since t h i s work was completed has shown that carbon monoxide reduces Ag(l) r a p i d l y i n basic solutions. A d ditional work has been done on t h i s system i n acid s o l u t i o n . Experimental Mate r i a l s The CO and C0-N 2 mixtures were reagent grade gases supplied by the Matheson Company, Inc., East Rutherford, New Jersey. The CO contained 35 ppm Hg and was used as received. Nitrogen used f o r e q u i l i b r a t i o n s was supplied by Canadian L i q u i d A i r Company. Mercury (II) perchlorate solutions were prepared by d i s s o l v i n g accurately weighed portions of reagent grade mercury (II) oxide i n a known amount of p e r c h l o r i c acid. The stock solutions contained a s l i g h t excess of 68. acid. Mercury (II) oxide from both Fisher Chemical Company and from Baker and Adamson Chemical Company was used and gave i d e n t i c a l r e s u l t s . Potassium permanganate stock solutions were prepared from reagent grade material. The solutions were boiled and, a f t e r standing f o r several hours, were f i l t e r e d through washed glass wool. Deuterium oxide was supplied by Atomic Energy of Canada Limited. I t was d i s t i l l e d from a l k a l i n e permanganate before use. The DCIO^ was pre-pared by Dr. J . F. Harrod by successive e q u i l i b r a t i o n s of DgO with HCIO^. Other solutions and chemicals were prepared as described i n Part I, A n a l y t i c a l The concentrations of mercury (II) solutions were checked periodi-c a l l y by t i t r a t i n g with ammonium thiocyanate using an i r o n ( I I I ) i n d i c a t o r . The ammonium thiocyanate was standardized with mercury (II) oxide. S i l v e r was determined g r a v i m e t r i c a l l y by p r e c i p i t a t i o n as s i l v e r c h l o r i d e . Procedure The general arrangement of the apparatus used i n both the mercury (II) and permanganate reductions i s shown i n Figure 10. This i s s i m i l a r to that used and described by Webster (117). The CO a f t e r passing through a flowmeter was l e d through the presaturator A containing a s o l u t i o n of the same concen-t r a t i o n s of HCIO^ and NaClO^ as the reactant s o l u t i o n . From the presaturator the CO entered the reactant s o l u t i o n contained i n the vessel B. E f f e c t i v e contact between the gas and s o l u t i o n was established by d i s p e r s i n g the gas 70. through the sintered disc C near the bottom of the reaction vessel. The gas stream then continued out through the tube D and was fed into the air inlet of a Mekker burner to dispose of the unreacted CO. Sampling of the reaction solution was done by closing tap E and allowing the gas pressure to force the solution out through tap F. Both the presaturator and the reaction vessel were immersed in a thermostatted water bath controlled to 0.1°. In the experiments with Hg(II), the concentration of Hg(l) formed was determined by pipetting a portion of the sample into a known volume of an Ig solution. (Usually about 5 ml of sample was collected and 2 ml was taken for determination of the Hg(l)). The unreacted Ig was titrated with a thio-sulfate solution standardized with KBrO^ .. Soluble starch or Fisher "Thiodyne"' was used to determine the end points. To ensure that the Hg(ll) solutions were stable N2 was passed through for a time comparable to the length of the run, usually 20 minutes, and samples were taken during this period. It was not necessary to know the exact concentration of the 1^  solution. The con-centration of Hg(ll) remaining in the solution at the time of sampling was calculated by means of the formula [Hg(Il)] = 0 - (A - V) N/S where 0 = original concentration of Hg(ll), M A = volume of SgO^  required to titrate a l l of the I 2 , ml V = volume of S^O^-2 required at time t, ml N = normality of SgO^2" S' = volume of sample, ml Time zero was taken as the time when bubbles of CO appeared in the reaction vessel. Samples were withdrawn at predetermined intervals and the sample time was that time at which tap E was closed. Thus time zero was quite 71. a r b i t r a r y , but the time i n t e r v a l between samples was known to within 5 seconds. In the f a s t e r runs the time i n t e r v a l between samples was 100 seconds. I n i t i a l l y there was some scatte r of t i t r a t i o n values which w a s 3 a t t r i -buted to r e s i d u a l CO continuing to react with H g ( l l ) . Thi3 was overcome by adding i c e to the Ig solutions and keeping the Ig - Hg(Il) solutions i n a r e f r i g e r a t o r u n t i l the end of the run (about 30 minutes) when they could be t i t r a t e d . The permanganate reductions i n acid and neutral solutions were per-formed i n the same apparatus using the same sampling technique. For the uncatalyzed r e a c t i o n i n acid solutions or neutral solutions with added e l e c t r o -l y t e , the suspended MnOg was e a s i l y removed by ce n t r i f u g i n g . A known volume of the c l e a r s o l u t i o n was pipetted into a potassium iodide s o l u t i o n and the l i b e r a t e d iodine was t i t r a t e d with sodium t h i o s u l f a t e . The permanganate con-centration was c a l c u l a t e d from the formula [MnO,"] = VN 4 5S where V = volume of SgO^ - s o l u t i o n , ml N = normality of SgO^ so l u t i o n 3" = volume of sample, ml With n e u t r a l solutions containing no added e l e c t r o l y t e the MnOg remained as a c o l l o i d uniformly d i s t r i b u t e d i n the so l u t i o n . In t h i s case a c i d i f i e d potassium iodide was.used and the t i t r a t i o n included both the un-mm reacted MnO^ and the c o l l o i d a l MnOg. Cal c u l a t i o n of MnO^ concentration was made by means of the expression 72. where V0 = volume of SgO^ solution before passing in CO, ml. In the Ag(l) and Hg(ll) catalyzed reductions of MnO^  by CO i t was found that MnOg formed remained in the solution as a uniformly dispersed colloidal suspension for a considerable period of time even with a perchloric acid concentration of one molar. Consequently, i t was possible to titrate the combined MnO^  and MnOg. In alkaline solutions the primary product in the CO reduction of MnO^  was found to be MnO^-. Titrations of the combined MnO^ ~ and MnO^ ~ were erratic with unsatisfactory end points. A better method of following these reactions proved to be the measurement of the CO uptake at constant pressure. The apparatus used for this purpose is shown in Figure 11. The reaction flask A which can be immersed into a thermostatted bath is connected by means of a flexible capillary tubing to the gas burette B in a constant temperature bath. The gas burette is connected to a simple gas handling system containing a manometer and connections leading to a vacuum pump and a gas cylinder. In operation a known amount of alkaline permanganate solution was pipetted into the reaction flask and degassed by repeated freezing and thawing under vacuum. The condenser C minimized loss of water from the solution. The degassed solution was then allowed to come to temperature equilibrium under a low pressure of CO, usually about 100 mm. CO was then admitted to the system and the pressure adjusted to the desired value. The reaction flask was attached to a mechanical shaker and this shaker was then started. Taps D and E were closed after a delay of about 10 seconds. It was established that there was no uptake of CO by the quiescent solution. The rate of solution of CO was very much faster than the rate of reaction and in i t i a l saturation of the solution was complete in less than 10 seconds. CO was admitted to the burette through a needle valve to keep the pressure con-Figure 11. Gas absorption apparatus 74. stant. The capillary manometer F served as a pressure indicator. As the reaction proceeded the mercury rose i n the tube G and i t s level was measured from time to time by means of a vertically mounted travelling microscope. The diameter of the tube G was known so a direct measure of the volume of CO used was obtained. Ah attempt was made to measure the s o l u b i l i t y of carbon monoxide using the gas absorption apparatus. Because of the small volumes involved the accuracy was no better than 20%. The published values f e l l i n this range so they were used throughout. The solubility data used are given i n Table XVI and have been adapted from Seidel (123). Results and Discussion Oxidation of CO by Hg(ll) Carbon monoxide i s oxidized by Hg(ll) according to the equation CO + 2Hg 2 + + H20 > C0 2 + H g 2 2 + + 2H + There was no indication of further reduction of the Hg(I). The reaction was found to be f i r s t order i n both CO and Hg(ll) and to obey the rate law --Jg9- =k[co][H g(ii)] Experimentally, the reactions were conducted under a constant CO pressure and the rate was followed by determining the concentration of Hg(Il). Log [*Hg(lI)J was plotted against time and the slope of the straight line gave the pseudo f i r s t order rate constant, k'» This i s related to the second order rate constant, k, by k = 2.303k' 2 [COl 75. Table XVI Solubility of CO in water Temperature CO' °C M/mm CO pressure x 10( 0 2..08 2 1.98 4 1.89 6 1.81 8 1.73 10 1.66 13 1.55 16 I.46 20 1.36 25 1.26 30 1.17 35 1.11 40 I.O4 45 0.99 50 0.95 60 0.87 76. The f i r s t order rate plots i n Figure 12 are for d i f f e r e n t concen-trations of Hg(II). In a l l cases good straight l i n e s were obtained even when the reaction was allowed to proceed to over 90$ completion. This indicates that there was no further reduction of Hg(l) and that the Hg(l) produced had no influence on the k i n e t i c s . The plots show an induction period of about one minute which i s probably the time necessary to saturate the solutions with CO. Tables XVII to XIX contain the res u l t s of experiments i n which the concentrations of reactants, HCIO^ and NaClO^ were Varied. The results of two experiments conducted i n DgO are also included. The f i r s t order character of the reaction with respect to CO i s shown i n Figure 13. Here the f i r s t order rate constant, k', i s plotted against the p a r t i a l pressure of CO. There was some var i a t i o n of the second order rate constant when the HCIO^ concentration was varied from 0.01 to 2.0 M. The io n i c strength was maintained constant at 2.0 M by the addition of NaClO^. The var i a t i o n was not large, but there was a d e f i n i t e minimum at a HCIO^ concentration of about 0.2 M. The i o n i c strength of the medium also had a small, but d e f i n i t e effect on the rate as indicated i n Table XIX. I t i s l i k e l y that the rate increase with added NaClO^ which i s p a r a l l e l f o r HCIO^ concentrations of 0.01 and 0.2 M i s due to an increased s o l u b i l i t y of CO. The increase as NaClO^ i s replaced by HCIO^ above 0.2 M HCIO^ i s also probably due to a s o l u b i l i t y effect. A simi l a r phenomenon has been encountered i n reactions of hydrogen (.119). The increase i n rate as the perchloric acid concentration i s re-duced below 0.2 M can be explained i n terms of a contribution by an hydro-lyzed mercuric species, S i l l e n (120) has shown that i n a pH range from about 78. Table XVII Effect of reactant concentrations on the rate of the CO - Hg(ll) reaction Temperature 40.. 0° HCIO^ Hg(ll) P c o k' k M M x 102 atm sec" 1xlQ 4 M^sec* 1 0.19 1.00 0.92 6.5 1.02 0.19 0.99 0.92 6.7 1.05 (a) 0.19 0.80 0.92 6.5 1.02 0.19 0.40 0.92 7.0 1.10 1.00 1.00 0.92 9.5 1.46 1.00 1.00 0.92 9.0 1.43 1.00 1.00 0.62 6.5 1.52 1.00 1.00 0.31 3.0 1.40 2.04 1.00 0.92 11.1 1.72 (b) 2.04 1.00 0.92 12.4 1.92 (c) 2.04 1.00 0.92 12.3 1.92 (d) 2.04 1.00 0.92 12.8 2.02 2.04 1.00 0.62 8.8 2.09 2.04 1.00 0.31 4.0 1.87 DC104 in D20 2.01 1.11 0.92 8.0 1.22 2.01 1.11 0.92 8.3 1.29 (a) Reaction 95$ complete (b) Flow rate 0.24 l/min (c) Flow rate 0.60 l/min (d) Flow rate 0.90 l/min Table XVIII E f f e c t of HCIO^ on the rate of the CO - H g ( l l ) r e a c t i o n Ionic strength maintained at 2.0M with NaClO^ Temperature 4O.O0, CO 0.92 atm Hg(II) HCIO^ NaClO^ k M x 10 2 M M M"1 s e c " 1 1.00 2.04 - 1.77 1.00 1.11 0.92 1.74 1.00 0.56 1.47 1.59 1.00 0.37 1.67 1.55 1.00 0.19 I.84 1.47 1.00 0.085 1.95 1.61 1.00 O.O48 1.99 1.65 1.00 0.011 2.03 1.80 Table XIX Effect of NaClO^ on the rate of the CO - Hg(ll) reaction Temperature 40.0°, CO 0.92iatm *g(II) HCIO^  NaClO^ k 4 x 102 M M M^sec" 1.00 0.19 1.05 1.00 0.19 0..28 1.13 1.00 0.19 0.56 1.22 1.00 0.19 1.11 1.26 1.00 0.011 «. 1.36 1.00 0.011 0.014 1.26 1.00 0.011 0.014. 1.39 1.00 0.011 0.028 1.37 1.00 0,011 0.056 1.35 1.00 0.011 0.083 1.31 1.00 0.011 0.111 1.38 1.00 0.011 0.128 1.35 1.00 0.011 0.278 1.43 1.00 0.011 0.456 1.49 1.00 0. Oil 0.479 1.44 1.00 0.011 1.67 1.73 1.00 0. Oil 2.03 1.81 81. 12.0 8.0 4.0 CO pressure atm Figure 13. E f f e c t of CO pressure on rate of CO - Hg(II) reaction at 40.0°. Hg(II) 1.00 x 10"2M; HC10. 2.04 M 82. 2 to 4 the concentration of HgOH+ cannot exceed 14$ of the t o t a l H g ( l l ) present.. Thus, i t would be d i f f i c u l t to e s t a b l i s h q u a n t i t a t i v e l y the exis-tence of a reaction path i n v o l v i n g HgOH+. The temperature at which the reaction was c a r r i e d out was varied over the range 25° to 54° at HCIO^ concentrations of 1 and 2 M with no added NaClO^. The values of the various a c t i v a t i o n parameters were found to be nearly the same f o r both solutions and are l i s t e d i n Table XX,. These were calculated from the data given i n Table XXI and the Arrhenius p l o t s shown i n Figure 14. Table XX A c t i v a t i o n parameters of the CO - Hg(II) re a c t i o n at 40,0° HC10 4 k E & A P * AH* A S " M M^sec"^ kcal/mole kcal/mole kcal/mole e.u. 1.0 1.44 15.4 18.1 14.8 -10.6 2.0 2.00 14.2 17.5 13.6 -12.5 The k i n e t i c s of the reac t i o n r e v e a l that one molecule of CO and 2+ one Hg ion are present i n the activated complex. The l a c k of a pronounced dependence of the rate on i o n i c strength i s consistent with a reaction be-tween an ion and a neutral molecule. The behaviour of the rate as the HC10 4 concentration i s varied i n d i c a t e s that the unhydrolyzed Hg2"1" i s the main re a c t i v e species. As previously noted, at low a c i d i t i e s HgOH+ may be involved, but t h i s would appear to provide an ad d i t i o n a l path rather than being the main re a c t i v e species. Also i t i s u n l i k e l y that HgC10 4 or s i m i l a r species are involved. Table XXI Effect of temperature on the CO - Hg(ll) reaction HC104 Hg(ll) p rco Temp k M M x l O 2 atm °C M"1 sec 1..00 1.00 0.97 25.0 0.42 1.00 1.00 0.95 31.0 0. 69 1.00 1.00 0.92 40.0 1.43 1.00 1.00 0.92 40.0 1.46 1..00 1.00 0.90 46.0 2.30 2.04 1. 00 0.97 26.0 0.65 2.04 1.00 0.97 26.0 0.65 2.04 1.00 0.95 33.0 1.19 2.04 1.00 0.95 33.0 1.18 2.04 1.00 0.91 40.0 2.02 2.04 1.00 Cv91 40.0 1.98 2.04 1.00 0.89 47.0 2.88 2.04 1.00 0.88 47.0 2.90 2.04 1.00 0.84 54.0 5.00 2.04 1.00 .0.84 54.0 5.32 3.1 3.2 3 103/T °A Figure 14. Arrhenius plots f o r CO - Hg(II) reaction. A. LOOM HC10 4 B. 2.04M HCIO^ Hg(II) 1.00 x 10" 2 M 85. Oxidation of CO by MnO^" In acid and neutral solutions permanganate was reduced to manganese dioxide. This r e a c t i o n can be represented by the f o l l o w i n g equations 3C0 + 2Mn0^" + 2H + > 3C0 2 + 2Mn02 + H^ O or 3C0 + 2Mn0^" + fi^O > 3C0 2 + 2Mn02 + 20H~ Experimentally the rate of permanganate reduction was measured. This rate can be expressed i n terms of carbon monoxide oxidation by — = "3/2 = k[C0irMn0^  J The f i r s t order experimental rate constant, k', i s re l a t e d to k by k = 3/2 2.303k' [CO] Tables XXII and XXIII contain the k i n e t i c data obtained i n acid and neutral s o l u t i o n . In general, these confirm that the re a c t i o n i s f i r s t order i n both carbon monoxide and permanganate. There were only minor v a r i a -tions with changes i n a c i d i t y or i o n i c strength. Adding a d d i t i o n a l manganese dioxide and varying the gas flow rate had no e f f e c t thus demonstrating the homogeneous character of the reaction . Some t y p i c a l r a t e pl o t s are shown i n Figure 15. The e f f e c t of the p a r t i a l pressure of CO i s shown i n Figure 16. The r e s u l t s obtained with a p a r t i a l pressure of carbon monoxide of 0.56 atm i s worthy of comment. The p a r t i a l pressure of carbon monoxide was varied through the use of N 2 - CO mixtures of various nominal compositions. These compositions were checked by gas chromatographic analysis and were 86, Table XXII Kinetic data for the oxidation of CO by Mn04 in acid and neutral solutions at 50.0° Mn04~ CO HC104 NaClO^ k' k M x 103 atm. M M sec" 1 x 10 4 M sec 1.94 0.87 0.37 mm 2} 9 1.61 3.36 0 .88 0.37 - 2 .9 1.61 6.72 0.87 0.37 - 2 .9 1.61 2.18 0 .88 0.11 _ 3.0 I . 6 4 (a) 2.18 0 .88 0.11 - 3.2 1.76 (b) 2.18 0 .88 0.11 - 3.0 1.62 (c) 2Z.1B 0 .86 0.11 - 3.1 1.70 (d) 2.18 0 .88 0.11 3.1 1 .66 2.02 O.64 0.11 - 2 .5 1.86 2.02 0 .56 0.11 - 2 .9 2.49 (e) 2.02 0.56 0.11 - 2.7 2.21 (e) 2.18 0.59 0.11 - 3.3 2.65 (e) 2.02 0.41 0.11 - 1 .8 2.06 2.02- 0.28 0.11 - 1.0 1 . 6 4 2.18 0.30 0.11 - 1.0 1.58 2.02 0.87 0.11 mm 3.0 1.64 2.02 0.86 0.11 0.28 2 .8 1 .56 2.02 0.86 0.11 0.58 2.7 1.48 2.02 0.87 0.11 0.89 2 .6 1 .45 2.02 0.85 0.26 0.75 2.7 1.50 2.02 0.85 0.52 0.49 2.9 1.64 2.02 0.85 0.78 0.23 3.0 1.67 25.02 0.87 1.00 - 3.2 1.75 2; 02 0.87 _ 2.1 1.13 2.02 0 .88 - - 2.1 1.16 2.02 0.87 - 1.00 1 .8 0.97 2.02 0.86 - 1.00 1 .8 1.00 (a) Flow rate 0.24 l/min (c) Flow rate 1.14 l/min (e)' See text (b) Flow rate 0..50 l/min (d) Contained Mn02 20 mg/100 ml Table XXIII E f f e c t of temperature on the rate of the oxidation of CO by MnO^ i n acid and neutral solutions MnO/ PC0 HC104 Temp k 4 x 10 3 atm M °C M"-'- sec""-1.68 0.96 0.092 25.0 0.298 1.68 0.96 0.092 28.0 0.342 1.68 0.94 0.092 33.0 0.542 1.68 0.92 0..092 39.1 0.763 1.68 0.90 0.092 43.8 1.10 1.68 0.88 0. 092 50.0 1.66 1.68 0.96 _ 28.0 0.212 1.68 0.94 - 33.0 0.316 1.68 0.92 - 39.1 0.514 1.68 0.90 - 43.8 0.740 1.68 0.88 - 50.0 1.10 88. 89. 90. generally within 5% of the nominal values. The gas mixture i n question gave r e s u l t s which were obviously out of l i n e , but which were reproducible. This i s even more puzzling as the same mixture was used i n the CO - H g ( l l ) reac-t i o n and gave normal r e s u l t s . The simplest explanation i s that the mixture contained a small amount of an impurity which was r e a c t i v e toward permangan-ate, but not toward H g ( l l ) . In basic solutions carbon monoxide reduces permanganate to man-ganate, i . e . CO + 2MnO^" + 4OH > C 0 3 2 " + 2Mn0^2" + 2R"20 Manganese dioxide i s produced e i t h e r by f u r t h e r slow reduction of the man-ganate CO + MnO^2" > C 0 3 2 " + Mn0 2 or by a slow disproportionation 3 M n 0 4 2 " + 2 ^ 0 > 2 M h 0 4 ~ + ^ 2 + 4°H" The production of manganese dioxide i s slow compared to the reduc-t i o n of permanganate to manganate. Figure 17 shows the rate of absorption of carbon monoxide by an a l k a l i n e s o l u t i o n of permanganate at 30.0°. The f i r s t order rate p l o t f o r t h i s data i s shown i n Figure 18. No extensive examination of the k i n e t i c s of the r e a c t i o n i n a l k a l i n e s o l u t i o n was made. Table XXIV l i s t s the complete data i n c l u d i n g the e f f e c t of temperature. The second order rate constant, k, i s r e l a t e d to the experi-mental constant, k', by v - 2.303 k'  k " 2 [CO] 800 1600 2400 T ime s ec Figure 18. Rate pl o t f or data of Figure 17. Table XXIV Kinetic data for the CO - MnO/ reaction in alkaline solution Temp °C MnO^ " M x l O 2 p c o atm NaOH M k ~1 -, ^ 4 sec xlO k M "^ sec 30.0 1.69 0.95 0.54 1.66 0.22 40. 0 1.69 0.92 0.54 2.62 0.41 40.0 1.35 0.92 0.43 2.44 0.39 40.0 1.69 0.61 0.54 1.75 0.42 44.2 1.69 0.91 0.54 4.3 0.71 50.0 1.69 0.88 0.54 4.9 0.89 57.0 1.69 0.83 0.54 6.9 1.41 -1 (a) Contained 0.02 M Ba(C10.) 94. In one experiment BatClO^Jg was added to p r e c i p i t a t e the manganate as i t was formed. Since t h i s had no e f f e c t i t may be concluded that other reac-tions i n v o l v i n g manganate are unimportant i n t h i s system. Arrhenius p l o t s f o r the CO - MnO^ reactions i n acid, neutral and basic solutions are shown i n Figures 19, 20 and 21. The calculated a c t i -vation parameters based on the oxidation of carbon monoxide at 50.0° are presented i n Table XXV, Included i n the table are the r e s u l t s obtained by Just and Kauko, These values were ca l c u l a t e d from t h e i r data f o r neutral permanganate at 15° and 25°, I t i s seen that the agreement i s excellent. Table XXV Ki n e t i c parameters f o r the CO - MnO," re a c t i o n at 50.0° A F * A H " AS* pH k.M*7 s e c " 1 E a. kcal/mole kcal/mole kcal/mole e.u. 1.0 1.66 13.5 18.7 12.8 • -18.1 7.0 1.10 14.7 18.9 14.0 -15.2 13.7 0.89 13.6 19.1 12.9 -19.0 7.0 1.16 (a) H.7 18.5 U.1 -14.7 (a) Calculated from the r e s u l t s of G. Just and Y. Kauko (79) There i s a s l i g h t increase i n rate as the solution passes from a l k a l i n e to weakly acid. I t i s u n l i k e l y that t h i s r e f l e c t s any s i g n i f i c a n t change i n mechanism, A s i m i l a r trend has been encountered i n the reduction of permanganate by molecular hydrogen (117). C a t a l y s i s of the CO - Mn04" reaction by Ag(I) and H g ( l l ) When the e f f e c t of added cations on the CO - MnO;- r e a c t i o n was 96. 97. 98. examined i t was found that both Ag(I) and H g ( l l ) ions had pronounced cata-l y t i c e f f e c t s . The reduction of permanganate was so f a s t that the reaction had to be followed at lower temperatures with very low l e v e l s of the c a t a l y s t ions. The ions of Cd(II), C u ( l l ) , F e ( l I I ) and T l ( l l l ) d i d not function as c a t a l y s t s . I n i t i a l l y , the catalyzed r e a c t i o n appeared to be very e r r a t i c and poor rate p l o t s were obtained. I t was then discovered that the trouble was l a r g e l y caused by the f a i l u r e of the manganese dioxide to separate. This was unexpected as the reaction solutions were one molar i n pe r c h l o r i c acid and no trouble had been experienced with the uncatalyzed reduction of perman-ganate i n acid solutions. Apparently the rate of reduction of permanganate i n the catalyzed r e a c t i o n i s much f a s t e r than the rate of coagulation of the manganese dioxide, r e s u l t i n g i n a c o l l o i d a l d i s persion. The samples were withdrawn and pipetted i n t o a potassium iodide s o l u t i o n and the l i b e r a t e d iodine was t i t r a t e d with sodium t h i o s u l f a t e . That i s to say, the product Mn(lV) was included i n the t i t r a t i o n . Quite good f i r s t order rate p l o t s were obtained up to about 50$ rea c t i o n when a p r e c i p i t a t e of manganese dioxide appeared. Most measurements were made at 13° and as the temperature was increased the p r e c i p i t a t e appeared at lower conversion. I t was necessary to take the complete set of samples i n any one experiment and then t i t r a t e . Somewhat better r e s u l t s were obtained i f the samples were added to the potassium iodide s o l u t i o n as they were taken rather than waiting u n t i l the end of the experiment. To confirm the stoichiometries of the catalyzed and uncatalyzed reactions the following experiment was conducted. The carbon monoxide was passed through a s o l u t i o n of sodium hydroxide to remove any carbon dioxide and then passed through 24-6 ml of a solu t i o n 6,72 x 10" 3 M i n KMhO, and 0,4 M 99. i n HGIO^. The e x i t gas was then passed into a carbonate free s o l u t i o n of sodium hydroxide. The permanganate s o l u t i o n was maintained at 50.0° and the gas bubbled through f o r 60 minutes. At the end of the time barium c h l o r i d e was added to the terminal hydroxide solution to p r e c i p i t a t e barium carbonate. This p r e c i p i t a t e was c o l l e c t e d and weighed. The permanganate concentration of the r e a c t i o n s o l u t i o n was determined. I t was found that 1.23 x 10" 3 moles of permanganate had been consumed and 1.74 x 10 moles of barium carbonate had been c o l l e c t e d . /According to the assumed stoichiometry I.84 x 10 ^ moles should have been produced. The experiment was repeated with the permanganate s o l u t i o n being 7.93 x 10 M i n AgClO^. The temperature was 13.0° and gas was passed f o r 4.0 minutes. The permanganate consumed was 1.29 x 10"*3 moles. The barium carbonate produced was 1.96 x 10 moles which compares with a calculated amount of 1,93 x 10 ^ moles. In another experiment carbon monoxide was passed through a s o l u t i o n 0.11 M HC10 4, 2.02 x 10" 3 M KMnO^ and 0,021 3M AgClO^. At 50.0° t h i s solution became colo u r l e s s with a brown p r e c i p i t a t e of manganese dioxide within 300 seconds. The colourless s o l u t i o n was found to contain 0.0213 M AgClO^. . The foregoing experiments demonstrate that the stoichiometries of the uncatalyzed and the s i l v e r catalyzed reduction of permanganate by carbon monoxide are the same. Also there i s no detectable l o s s of s i l v e r i n the catalyzed reac t i o n . The mercury (II) catalyzed reaction has been assumed to proceed i n the same manner as the s i l v e r catalyzed r e a c t i o n . In Figure 22 rate plots f o r the permanganate reduction with and without c a t a l y s t are compared. The slopes of the f i r s t order plots y i e l d e d a pseudo f i r s t order rate constant which included both the catalyzed and 800 1600 2400 T ime s ec Figure 22. Rate plo t s for uncatalyzed and catalyzed CO - MnO^ reaction at 13.0°. A. no c a t a l y s t -4 B. Hg(Il) . 4.00 x 10 M C. Ag(I) 7.93 x 10"6M 101. uncatalyzed reactions. The c o n t r i b u t i o n of the uncatalyzed reaction was sub-tracted before values of the t h i r d order rate constants were calculated. The e f f e c t of the reduction of Hg(II) and the reoxidation by permanganate i s n e g l i g i b l e at the low temperatures and low H g ( l l ) concentrations employed. The k i n e t i c data f o r the H g ( l l ) catalyzed reaction are given i n Table XXVI. Corresponding data f o r the Ag(l) catalyzed r e a c t i o n are given i n Table XXVII. The f i r s t order rate constant, k', has been corrected by subtracting the c o n t r i b u t i o n of the uncatalyzed reaction. The t h i r d order rate constant of the catalyzed reaction, k , i s obtained from k' by k = 3/2 2.303 k' k 3/2 mm~ where [M] i s the concentration of H g ( l l ) or A g ( l ) . While the v a r i a t i o n i n the c a l c u l a t e d rate constants i s about 20$ i t seems safe to conclude that the catalyzed reaction i s f i r s t order i n each of permanganate, carbon monoxide and metal ion. There i s a small apparent decrease i n the rate constant of the Ag(l) catalyzed reaction as the permanganate concentration increases. The explanation of t h i s e f f e c t , i f r e a l , i s not c l e a r . The e f f e c t of temperature on the rates of the catalyzed reactions i s shown by the data given i n Tables XXVIII and XXIX. In the temperature range 0 - 25° the r a t e constants increased by a f a c t o r of 3 f o r the H g ( l l ) catalyzed r e a c t i o n , but only by 50$ f o r the Ag(l) catalyzed reaction. Ar-rhenius plots were constructed from the Series B data of the Tables and are shown i n Figures 23 and 24* The a c t i v a t i o n parameters f o r 0° are given i n Table XXX. Table XXXVI Kinetic data for the Hg(ll) catalyzed oxidation of CO by MnO Temperature 13.0° HC10, 1.00 M MnO/ Hg(H) CO i k- k i A x 103 M x IO 4 M x IO3 sec"1 x IO 4 M^sec"1 : 3.4 4.00 1.15 2.8 2.1 6.7 4.00 1.15 2.6 2.0 10.1 4.00 1.15 2.3 1.7 13.4 4.00 1.15 2.4 1.8 16.8 4.00 1.15 2.3 1.7 20.2 4.00 1.15 2.3 1.8 10.1 2.00 1.16 1.1 1.7 10.1 8.00 1.16 4. 4 1.6 6.7 4.00 0.40 0.7 1.8 6.7 4.00 0.53 1.3 2.2 6.7 4.00 0.'82 2.1 2.2 103. Table XXVII Kinetic data for the Ag(l) catalyzed oxidation of CO by MnO^ " Temperature 13.0° HCIO^ 1.00 M Mn04" Ag(I) CO k' k c M x 103 M x 106 M x IO3 sec"1 x IO 4 M^sec"1 x IO"5 2.0 7.93 1.15 5.2 2.0 3.4 7.93 1.15 4, 4 1.7 6.7 7.93 1.16 3.7 1.4 10.1 7.93 l . H 3.6 1.4 13.4 7.93 1.14 3.4 1.3 6.7 3.96 1.15 2.1 1.5 6.7 15.9 l . U 6.1 1.2 6.7 31.7 1.16 11.6 1.1 6,7 7.93 0.40 1.2 1.3 6.7 7.93 0..53 1.9 1.5 6.7 7.93 0,.82i 2>,8 1.5 104. Table XXVIII Series A Series B Effect of temperature on the rate of the Hg(ll) catalyzed CO - MhO^ - reaction MnO " 4 CO Hg(ll) Temp k c M x IO3 M x IO 4 M x IO 4 °C M~2sec"*1 x 6.72 9.2 4.00 25..0 3.58 6.72> 10.0 4.00 20.0 3.12 6.72 10.8 4.00 16.0 2.44 6.72 12.4 4.00 10.2 1.87 6.72 15.5 4.00 0.1 1.27 7.-88 9.43 4.00 2 5.0 3.26 7.88 10.2 4.00 20.4 2.66 7.88 10.9 4.00 16.7 2.23 7.88 11.8 4.00 13.0 1.85 7.88 12.6 4.00 9.1 1.61 7.88 , 15.6 4.00 0.1 1.09 ,-3 Table XXIX Effect of temperature on the rate of the A'g(l) catalyzed CO - MnO"^ " reaction MnO^ " CO Afc(I) Temp kc M x 103 M x 10 4 M x IO 6 °C ur2 -1 M sec x Series A 6.72 9.2 7.93 25.0 1.83 6.72 10.0 7.93 20.0 1.55 6.72 10..8 7.93 16.0 1.57 6.72 12.4 7.93 10.0 1.43 6.72 15.5 7.93 0.1 1.29 Series B 7.88 9.4 7.93 ' 25.0 1.47 7.88 10.1 7.93 21.0 1.47 7.88 10..9 7.-93 16.7 1.30 7.88 11.8 7.93 13.0 1.28 7.88 15.6 7.93 0.1 1.10 3.4 3.5 3.6 •103/T °A Figure 23. Arrhenius p l o t of Hg(II) c a t a l y z e d CO - MnO^ r e a c t i o n . Hg(II) 4.00 x 10 _ 4M; CO 0.98 atm; MnO " 7.88 x 10"3M; HC10, LOOM 5.2 3.4 3.5 3.6 103/T °A Figure 24. Arrhenius plot of Ag(I) catalyzed CO - MnO^ reac t i o n Ag(I) 7.93 x 10"6M; CO 0.98 atm; MnO." 7.88 x 10 _ 3M; HC10, 1.00 M 108. Table X X X A c t i v a t i o n parameters f o r the catalyzed CO - MnO," reactions y o ° ) E act AF* A H * AS* M ^ s e c " 1 kcal/mole kcal/mole kcal/mole e.u. Hg(H) 1.09 x 10 3 7.0 12.2 6.5 -20.7 Ag(D 1.10 x 10 5 1.9 9.6 1.3 -30.4 109. Tracer Studies with 0 X O The general procedure was s i m i l a r to that used by Wiberg and Stewart (121) f o r the formate-permanganate r e a c t i o n . In o u t l i n e , the pro-cedure was to pass carbon monoxide through an a l k a l i n e s o l u t i o n of the l a b e l l e d permanganate, destroy the permanganate and manganate with hydrazine, remove the manganese dioxide and p r e c i p i t a t e the product carbonate as barium carbonate. The barium carbonate was decomposed and the evolved carbon dioxide was c o l l e c t e d i n the apparatus- shown i n Figure 25. The p r e c i p i t a t e of barium carbonate was placed onto frozen concentrated s u l f u r i c acid con-tained i n the small lower bulb and the apparatus evacuated. The s u l f u r i c acid was melted by immersing the bulb i n warm water whereupon carbon dioxide was released and c o l l e c t e d i n the large upper bulb. The l a b e l l e d permanganate contained 1.14$ and was supplied by Dr. J . B. Farmer. The mass spectrometric analyses were performed by Dr. D. C. Frost. In these experiments i t was convenient to compare mass 46 ( C 1 2 0 l 6 0 1 8 p t o m a s g 4 5 ( C 1 3 a l 6 ^ , g g t h e g e a r e o f t h Q g m e o r d e r o f magnitude and the value of mass 45 remains e s s e n t i a l l y constant f o r COg obtained from 18 natural sources and 0 enriched sources. The natural abundances f o r the isotopes of carbon and oxygen are (122) i n per cent C 1 2 98..892, C 1 3 1.103 , 0 1 6 99.758 , 0 1 7 0.0373 , 0 1 8 0.2039. This leads to a r a t i o of mass 46 to mass 45 equal to 0.343. The experimen-t a l l y determined r a t i o using C0 2 formed by o x i d i z i n g CO with unlabelled Mn04" was 0,332 i 0.027. The r e s u l t s obtained with l a b e l l e d permanganate follow. Figure 2 5 . Apparatus for decomposing BaCOo and c o l l e c t i n g CO 111. Test 1. 0.999 g KMhO^ i n 25 ml of 0.2M NaOH at 25°. 47$ CO was bubbled 2-through f o r 60 minutes. The MnO^ and MnO^ were destroyed with NgH^ and MnOg removed by centr i f u g i n g , BaClg was added and the r e s u l t i n g p r e c i p i t a t e of BaCO^ was washed and drie d . The p r e c i -p i t a t e was decomposed on s u l f u r i c acid and the evolved COg c o l -lected.. The mass 46/4-5 r a t i o was 0.409 which corresponds to 0.10 atom of oxygen transferred from permanganate per CO, Test 2. 0.099g KMnO^ i n 25 ml water at 25°. 47$ CO bubbled through f o r 40 minutes. The same procedure was followed. The mass 46/45 r a t i o was 0,352. E s s e n t i a l l y no oxygen was transferred from the permanganate. Test 3. 0.448g KMnO^ i n 50 ml 0.7M NaOH at 25°. Bubbled 100$ CO through the s o l u t i o n and took 10 ml samples after 5> 20, 45 and 80 minutes. The BaCO^ p r e c i p i t a t e s were c o l l e c t e d and treated as above. The COg samples were inadvertently d i l u t e d with a i r and only the 45 and 80 minute samples could be analyzed. Both gave 18 r a t i o s of 0,32 i n d i c a t i n g no 0 tra n s f e r . Test 4. 0.032g KMnO, i n 25 ml of 0.2 M Ba(0H)_ at 25°. 100$ CO was passed A- 2 through f o r 80 minutes. The combined p r e c i p i t a t e s of BaMnO^ and BaCO^ were c o l l e c t e d and the BaMnO^ decomposed with hydrazine. The BaCO^ was not separated from the MnOg before decomposition. The mass 46/45 r a t i o was 0.542 which i s equivalent to an 0 1 8 t r a n s f e r of 0.30 atom per CO. Test 5. 0.114g KMnO^ i n 50 ml of 0.2M Ba(0H) 2 was reduced with 1 ml of a 10$ formic acid s o l u t i o n . The pr e c i p i t a t e s were treated as i n the 18 previous t e s t . The mass r a t i o was 0,42 or 0,12 0 atom was trans-f e r r e d per CO. 112. Test 6. 0.078g KMnO i n 30 ml l.OM NaOH was reduced with 1 ml of 10$ 4 formic acid. The manganate was reduced with hydrazine and the MnOg removed, BaCO^ was pr e c i p i t a t e d by adding BaCL, and was treated i n the usual manner. The mass r a t i o was 0,507 correspon-18 ding to 0.25 0 atom transferred per CO. Test 7. 0.098g KMnO. i n 50 ml of a solu t i o n 0.6M i n NaOH and 0.08M i n 4 BaClg. 1 ml of a 10$ formic acid s o l u t i o n was added. The r e -s u l t i n g BaMnO^ was converted to MnOg ^-^ n hydrazine and the com-bined BaCO^ and MhOg Pr e c i p i t a t e s were c o l l e c t e d . The evolved C0 2 had a 46/45 mass r a t i o of 0.479. 0.21 0 1 8 atom was transferred per CO, Test 8. O.lOOg KMnO^ i n 25 ml of 1.8M NaOH at 4 O 0 . CO was passed through f o r 60 minutes and the carbonate c o l l e c t e d as usual. The mass r a t i o was-0,410 i n d i c a t i n g a tr a n s f e r of 0.10 atom of O 1^ per CO. Test 9. This t e s t was done pr i m a r i l y to check the stoichiometry of the s i l v e r catalyzed reaction, 0,140g of l a b e l l e d KMnO^ was dissolved i n 50 ml of a so l u t i o n 0,9M i n HCIO^ and 1.6 x 10"4 M i n Ag +. CO was passed through f o r 40 minutes at 40°. The e x i t gas was l e d through a solu t i o n of NaOH. BaClg was added to p r e c i p i t a t e BaCO^ which was c o l l e c t e d , d r i e d and weighed. The p r e c i p i t a t e was divided into two portions f o r generation of CO,,. Both gave mass 46/45 r a t i o s of 0.30 i n d i c a t i n g no 0 transfer. Test 10. 0.l60g of KMn04 i n 70 ml of a so l u t i o n l.OM i n NaOH and 0.06M i n BaClg. CO was bubbled through f o r 40 minutes at 40° and the com-bined p r e c i p i t a t e s treated as usual. The mass r a t i o was found to 18 be 0.763 which i s equivalent to an 0 atom tr a n s f e r of 0,64 per CO. The tests are summarized in Table XXXI. 113. The results show no clear pattern. The results obtained for the formate reduction of permanganate are in fair agreement with those of Wiberg and Stewart (121), These authors showed that the method did not involve any exchange of carbonate in basic solution. Mechanism of the CO - Hg(ll) Reaction A" satisfactory mechanism must account for the kinetic behaviour including the lack of a pH dependence. It must also provide for the stoi-chiometry, A mechanism which does this is the followingJ 1. - Hg 2. - Hg 3. The rate determining step must involve the aquo ion rather than + HgOH because the reaction i s essentially acid independent. The lack of acid inhibition also indicates that the reversal of the rate determining step cannot be important. This also shows that the formation of the intermediate is not a rapid equilibrium with the rate determining step being its decom-position. The subsequent fast reaction between Hg(ll) and Hg(0) is postulated for other reactions of mercury (123). The unique feature of this mechanism is the insertion of a molecule of CO between an Hg2* ion and a coordinated water molecule. The subsequent 0 2+ k " + + - 0H2 + CO » - Hg - C - OH + H rate determining 0 it - C - OH > Hg + C02 + H + fast Hg + Hg 2 + > Hg22+ fast 114. Table XXXI Summary of 0 transfer experiments Test 5 7 6 Reductant HCOO" n n< Atoms 0 from MnO^  0.12 0.2 0.25 T°C 20 20 20 13.3 13.8 14.0 Ba M 0.2 0.08 9 2 3 1 4 8 10 CO 0.0 0.0 0.0 0.10 0.30 0.10 0.64 40 25 25 25 25 40 40 0.0 7.0 13.8 13.3 13.3 14.3 14.0 0.2 0.06 115. decomposition of the intermediate takes place by loss of a proton leading to the products. This insertion of a CO between a metal-oxygen bond is somewhat analogous to the mechanism proposed by Sternberg and Wender (100) to explain the Co2(C0)g catalyzed formation of methyl formate according to CH 30H + Co2(C0)g CH-OCo(CO). + CO i 4 CH30C0Co(C0)4 + Hg The last step may be replaced by CH-OCOCo(CO). + HCo(CO). .7 4 4 -> CH30Co(C0)4 + HCo(CO)^ CH.OCOCo(CO). 3 4 -> HC00CH3 + HCo(C0)4 -> HC00CH3 + Co2(C0)g Particularly strong support for the proposed insertion mechanism of the CO - Hg(ll) reaction is given by the following consideration. Replace-ment of the proton of the intermediate by an alkyl group should lead to stabilization and allow the isolation of the intermediate. This has been shown to be the case. Schoeller et al (124) observed that methanolic solutions of Hg(ll) acetate readily absorb CO to form a compound with the formula HgC^ H^ O^ . The reaction is reversible, CO being liberated by heating or by treatment with acid. Schoeller proposed the structure 0 II 0 it CH3 - C - 0 - Hg - C - 0CH3 corresponding to the insertion of a molecule of CO between a monoacetate Hg(ll) and a methoxy group. On the other hand, Manchot (125) proposed a 116. carbonyl structure 0 ti CH„ - C 3 - 0 - Hg - OCH3 C 0 Halpern and Ket t l e (126) recently resolved the structure of the compound and showed that i t was that proposed by Schoeller. As noted previously, other CO i n s e r t i o n reactions are known. I t i s not yet known whether these are formed d i r e c t l y or whether a carbonyl structure i s f i r s t formed which then rearranges to give the i n s e r t i o n compound. The l a t t e r s i t u a t i o n seems to p r e v a i l i n the formation of acetylmanganese pentacarbonyl. The f a c t that the reduction of H g ( l l ) by CO i s slower i n D 20 than i n HgO can be interpreted as supporting a mechanism i n which a water molecule i s displaced. The r a t i o of the rate constants k / l c i s 1,6 at 40°, HgO ^2 Such a change i n rate constants can be accounted f o r by an increase i n a c t i -vation energy of about 0.3 kcal/mole. This of the same order as the d i f f e r -ence i n hydration energies of some ions i n D 20 and H^O, although an unambig-uous i n t e r p r e t a t i o n i s not yet possible (127, 128, 1 2 9 ) . Kemp (130) has examined the k i n e t i c s of the reaction 0 rti C l - Hg - C - OCH^ + H + C l ~ > HgClg + CO + CH^OH i n methanol. The rate was found to be f i r s t order i n hydrogen ion concen-t r a t i o n . This suggests that the reverse reaction leading to the formation of the carbon monoxide i n s e r t i o n compound involves CH^OH rather than CH^O . By analogy, the intermediate i n aqueous sol u t i o n involves H 20 rather than OH 117. Results of Related GO Reactions Following completion and publication of the present work (131) other CO reactions have been examined and the results support the concept of an insertion mechanism. McAndrews (132) has examined the kinetics of the Ag(I) - CO reaction in perchlorate and acetate media. Experimentally this is a d i f f i -cult system. The reactions are slow and i t was necessary to perform the experiments in an autoclave employing pressures up to 30 atm. and temperatures of 60 - 110 . The rate of reaction was followed by means of the pressure changes. The experimental rate law was = ka[C0][AgIc] + VcLcoJUg+jW]"1 -1 + kcKc [CO] [ Ag+] [ AgAc] [ H +J where Ac represents the acetate ion. The first term is acid independent and presumably reflects direct participation of AgAc in the reaction. The second term is independent of acetate concentration and is the same as Bauch's expression (116). The third term is complex. The mechanism proposed to account for the rate law is Ag + Ac N AgAc rapid equilibrium 0 Ka AgAc + CO > Ag - C - Ac 0 it Ag - C - Ac + Ag + + H20 > 2Ag + COg + HAc + H + fast 118. 0 K Ag + CO + H O Ag - C - OH + H rapid equilibrium 2 0 II Jr Ag - C - OH + 4 g + — 2 A g + C0 2 + H + 0 " k Ag - C - OH + AgAc — 2 A g + C0 2 + HAc In a complex system such as t h i s i t i s d i f f i c u l t , i f not impossible, to a r r i v e at d e f i n i t e conclusions about mechanism. I t i s of i n t e r e s t , how-ever, that two carbon monoxide i n s e r t i o n complexes are postulated. Nakamura and Halpern (133, 134.) have found that Ag(l) amine com-plexes react r a p i d l y with carbon monoxide to p r e c i p i t a t e m e t a l l i c s i l v e r . The k i n e t i c s are i n accord with the following mechanisms + J L , + AgL 2 + HO L-Ag-OH + LH rapid equilibrium 0 II L-Ag-OH + CO L-Ag-C-OH k - l 0 in L-Ag-C-OH + Ag(l) — p r o d u c t s L represents the amine and K = K ^ I ^ which are defined by d [ W ] K, = [HL +] [0H-] T> [ L J K h = CAg-L-OH1 [AgL +] [OH" ] 119. The experimental rate constant, k, i s defined by the rate law -1 . = Tclcol T A P L „ + 7 / * H L + ' dt - d C M = k [ C 0][AgL 2 +jrHL +jwhich i s equivalent to -<lM2 = k l [ C 0 J [L-Ag-OH] dt That i s , k = k K K K 1 d b h the important feature i s that the variation i n k i s very nearly accounted for by variation i n K^K^ so that k^K^ i s insensitive to the nature of the amine. This means that the increase i n pH accounts f o r the increase i n rate on going from acidic to amine-buffered media. For primary amines, the insertion of CO i n the L-Ag-OH complex was found to be the rate determining step and the above rate law describes the behaviour. With triethylamine, the reverse reaction, k_-^ , becomes fa s t enough to compete with kg and the Ag(l) species i s i d e n t i f i e d as AgL , When ammonia i s used with high concentrations of NH^4" i t appears that the Ag(l) species i s L-Ag-OH and the reaction becomes second order i n Ag(l) and + inverse second order i n NH^ . Byerley and Peters (135) have reported on the CO reduction of Cu(Il) to Cu(l) i n aqueous solutions at quite high CO pressures. The rate law can be expressed as the sum of two terms, _ -1 2 - 1 " = ^ [ C u d l ) ] [ Cu(l)] [ H +] + kgKg [Cu(ll)] [CO] [ H +J The mechanisms which are postulated to give t h i s rate law are 120 Path 1. Cu(CO) + HgO 0 it Cu - C - OH + Cu(ll) CuH+ + Cu(ll) Path 2. Cu(Il) + 0 0 + ^ 0 0 it - Cu - C - OH+ + Cu(ll) 2Cu(l) + 2C0 high pressures All of the foregoing reactions are postulated to proceed via CO insertion mechanisms. One way in which they differ from the corresponding Hg(II) reaction is that the decomposition of the insertion complex must occur by reaction with another metal ion. This is apparently slow enough in many cases so that the reversal of the intermediate forming step becomes important. Mechanisms of the Uncatalyzed and Catalyzed CO - MnO^ - Reactions The reactions between CO and MnO^  in basic, neutral and acidic solutions exhibit the same kinetics and have very similar activation para-meters. Consequently, i t is reasonable to assume that the rate controlling step is the same in each case. The different stoichiometries reflect the differences in stability of various manganese species with respect to pH, The proposed mechanism is the reduction of Mn(VIl) to Mn(V) by CO taking place in two steps: 0 ==± Cu - C - OH + H k l + -^4 Cu(l) + CuH + CQ fast at 2 ^ 4 2Cu(l) + H + 0 = ^ - Cu - C - OH + H > 2Cu(l) + C02 + H -> 2Cu(C0) + 121. 1. MnO/ + CO — 0 3 M n O - CO" 2. 03MnO - CO" + . HgO f a S t ) Mno/" + COg + 2H + 3-The hypomanganate, MnO^ , then reacts r a p i d l y to give the f i n a l manganese species. In basic s o l u t i o n the reaction i s probably 3- 2-MnO, + MnO, > 2MnO, 4 4 4 which i s known to be f a s t (136). In acid solution both manganate and hypo-manganate undergo rapid disproportionation to permanganate and manganese dioxide. An a l t e r n a t i v e formulation of the rate determining step i s CO + MnO/ —^-> C0 2 + Mn03" However, since Lux (137) has shown that Mn(V) s a l t s are isomorphous with 3-phosphates the formulation MnO^ i s preferred. The simplest method of depicting the rate determining step i s by a n u c l e o p h i l i c attack by a permanganate oxygen on the carbon atomoof carbon monoxide. This intermediate i s then decomposed by water which can attack e i t h e r the carbon atom or the manganese atom, i . e . OC 0 - - MnO, i J i 0 H2 or OC 0 - - MnO.," i ' i i 0 H 2 122. In the f i r s t case oxygen i s not transferred from MnO^~ to CO, whereas i n the second case i t i s . Water has been shown as the attacking agent, but the hydroxyl ion would produce the same r e s u l t , A mechanism s i m i l a r to that envisioned f o r the H g ( l l ) - CO reac-t i o n i n v o l v i n g a carbon monoxide i n s e r t i o n complex seems u n l i k e l y i n t h i s case. Carbon monoxide i s a poor n u c l e o p h i l i c reagent and requires consider-able s t a b i l i z a t i o n by back donation of electrons to i t s empty p i o r b i t a l s . The formation of metal carbonyls i s at t r i b u t e d to t h i s s t a b i l i z i n g e f f e c t because those metals which form carbonyls have d electrons available, Metals with high oxidation states such as Mn(VTI) have no electrons a v a i l a b l e f o r back donation so s t a b i l i z a t i o n of a CO complex i s not possible. The oxidation of CO by a l k a l i n e permanganate i s analogous to the oxidation of the i s o e l e c t r o n i c cyanide ion. This l a t t e r r e a c t i o n has been studied by Freund (138) and by Stewart and Van der Linden (139). The k i n e t i c s are complicated apparently because at low pH HCN takes part and at very high pH a mechanism invo l v i n g hydroxyl r a d i c a l may be important. At a pH between 13 and 14 the dominant r e a c t i o n path i s described by the rate law - dj£Cl = k[ C N-][MnOr] dt *^ which has the same form as that f o r the CO r e a c t i o n . Stewart and Van der Linden (139) found that oxygen tr a n s f e r from MnO^ to CN" occured and the proportion transferred increased with in c r e a s i n g pH. This may i n d i c a t e that i n the decomposition of the proposed intermediate NC 0 - - Mn0 3 2" both HgO and OH are e f f e c t i v e with the OH" attack occurring p r e f e r e n t i a l l y 123. on manganese. Since attack on the manganese lends to oxygen transfer this accounts for the increased transfer with increased pH. & similar situation may well prevail in the corresponding CO reaction. In the catalyzed CO - MnO^  reaction there is no reason to consider that the catalytic actions of Ag(l) and Hg(ll) are any different. Accordingly they have both been assumed to act in the same way. In view of the likelihood of a CO insertion complex providing the mechanism for the CO - Hg(ll) reaction i t is attractive to postulate a similar complex for the present reaction. When this is coupled with the proposal for the uncatalyzed reaction, the postulated complex becomes 0 it - Ag - C - OMnO^  The mechanism is then written as 0 k " Ag + CO + Mn04" ^ Ag - C - 0Mn03 rate determining 0 ti Ag - C - OMnO, + H-0 > Ag+ + CO, + MnO,3" + 2H+ 2 4 f a s t A"s with the uncatalyzed reaction the hypomanganate reacts rapidly to give Mn02. The catalytic action of Ag(l) and Hg(ll) must be to stabilize the equivalent of an insertion complex. This may be accomplished by donation of metal d electrons to form a pi bond with the CO molecule. Further comments along these lines will be made later. Comparison of CO, Hg and HCOOH as Reductants 124. A comparison of some of the reactions of hydrogen and formic acid with those of carbon monoxide i s of i n t e r e s t . A summary of k i n e t i c data f o r various reactions i s given i n Table XXXII. Halpern (1, 145) has discussed the reactions of hydrogen and Taylor and Halpern (77) have discussed those of formic acid. When hydrogen i s oxidized the molecule i s s p l i t . This s p l i t t i n g can take place e i t h e r homolytically to give hydrogen atoms or h e t e r o l y t i c a l l y to give an hydride ion and a proton. Both types of behaviour seem l i k e l y , but the same form of intermediate can be postulated i n each case. For ex-ample, C u ( l l ) s p l i t s hydrogen h e t e r o l y t i c a l l y according to the mechanism C u 2 + + Hg > CuH + + H + I t i s l i k e l y that Ag(l) i n the bimolecular r e a c t i o n , Hg(Il) and MnO^- act i n the same manner. On the other hand Hg(l), &g(l) i n the termolecular reaction and i n the catalyzed MnO^ reaction s p l i t hydrogen homolytically H g g 2 + + Hg > 2HgH + An examination of the electron configuration of the active cations shows that they have nearly f i l l e d or j u s t f i l l e d d - o r b i t a l s . The hydrogen electrons can then be accommodated i n the d - o r b i t a l s which are vacant or can be made vacant by promotion to the next s - o r b i t a l . Thus, i t can be seen that one of the f a c t o r s i n the s t a b i l i z a t i o n of the hydride intermediates i s the energy separation of the d- and s- l e v e l s . For a f i x e d quantum number, the energy separation depends on the nuclear charge of the ion. This explains 2+ o + why Hg activates hydrogen while the i s o e l e c t r o n i c ion Tlr i s i n a c t i v e . As the quantum number increases the d-s energy separation decreases. Thus the 125. Table XXXII Summary of kineti c data for some oxidations of H2, HGOOH , HGOO and CO Reactants Rate Law A H * Kcal AS* e,t,Pt Temp °C Ref H2-Cu2+ krH 2 ][cu 2 +j 26 -10 80- HO HO -Hg 2 + k[H 2][Hg 2 +J 18 -12 65-100 123 n 2 + -Hg2 k[H2][Hg 2 2 +J 20 -10 65-100 123 -Ag+ k ^ A g 4 " ] H -25 30-120 141 -Ag+ k[H 2]tAg +] 18 -12 30-120 141 -MnO^ " k[H2][MnOA_] H -17 30-70 142 -MnO^ " + Ag + k[H2][Ag+J[MnO^J 9 -26 30-60 142 HCOOH-Cb(Hl) k[HC00H][co(IIl)J 26 19 0-30 143 -MnO^ " k[HCOOH]fMnO^ " ] 16 -19 15-35 144 HCOCf-Co(lIl) k[HGOO"] [Co(lll)] 21 21 0-30 143 -Hg 2 + k[HC00-][Hg2"f] 20 3 36-61 77 u 2+ -Hg2 k [HCOO-J [kg/ 4] 21 0 60-80 77 -Tl3 + k[HC00-][Tl 3 +] 26 ~ 21 65-85 77 -MnO^ " k[HCOO-][MnO^ "] 12 - H 15-35 144 -MnO^ "+ Fe 3 + k [HCOO"] [FeMn0^2+J - - 30 144 2+ CO-Hg^  k H ^ ] 15 -13 26-54 -Ag+ k[CO][Ag+J [ H + ] ~ 17 - 80-110 132 -AgL2+ (a) k[C0][LAg0H] 9 -15 15-35 134--Cu(ll) k J c u d i O ^ c u d ) ] ^ ] +k 2[C0][cu(H))[H +J - - 120 135 -MnO^ k [C0][MhO 4 "] 13 -17 28-50 -MnO^ "+ Hg 2 + k[C0j[Mn0^']fHg2+J 6 -21 0-25 -MnO^ ~+ Ag + k[co][MnOA-][Ag+J 1 -31 0-25 (a) L = methylamine 126. 5 d 1 0 i o n H g 2 + i s active while the corresponding 4d and 3d ions C d 2 + and Zh 2 are i n a c t i v e . The permanganate reaction i s somewhat d i f f e r e n t , but may also 2-involve hydride ion t r a n s f e r . In t h i s respect the i s o e l e c t r o n i c CrO^ ion i s i n a c t i v e . This may be at t r i b u t e d to the l a r g e r energy gap between the 2-highest occupied and lowest unoccupied o r b i t a l s of CrO as compared to 4 MhO " (146). I t appears that A g + catalyzes the reaction between IL and MnO 2 4 by allowing a one electron t r a n s f e r (H atom) to MnO^ , i . e . Ag + + MnO." + H > A'gH+ + HMnO " 4 2 4 The oxidation of formic acid takes place by two paths one in v o l v i n g the molecule HCOOH and the other a t t r i b u t e d to the formate ion HCOO~. 2+ 2+ 3+ The ions Hg 2 , Hg and T l apparently react with HCOO by the same mechanism which involves coordination with the oxygen of formate i o n . This permits the t r a n s f e r of two electrons to the metal ions. The product T l + i s formed d i r e c t l y . With the mercury cations mercury atoms are the d i r e c t products of the reaction. The mechanism of the C o ( l I I ) reaction may be d i f f e r e n t . C o ( l l l ) i s a powerful one electron oxidant and both the formate ion and formic acid reactions may take place by an hydrogen atom t r a n s f e r . The permanganate oxidations are quite d i f f e r e n t f o r the two paths. The formate ion r e a c t i o n quite l i k e l y takes place by a hydride ion t r a n s f e r to MnO/. This i s i n l i n e with Wiberg and Stewart's (121)' observation that there i s a large deuterium isotope e f f e c t . The reaction i s not altogether that simple, however, as there i s also considerable oxygen t r a n s f e r from 12' MnO^ ~ to the product COg. There is no deuterium isotope effect in the formic acid reaction and i t has been postulated that the mechanism involves nucleophilic attack on the carbon atom by a permanganate oxygen, OH i ~0 - C - 0 - MnOo i H 3+ The catalysis by Fe of the oxidation of formate by permanganate is attributed to a mediating effect permitting the formation of Mn(Vl) rather than Mn(V) as the i n i t i a l reduction product according to HGOO" + MnO " + Fe 3 + > H + + GO + MnO.2' + Fe 2 + 4 <s 4 The catalysis has been assumed to apply to the formate reaction rather than the formic acid reaction although this has not been definitely established.. The carbon monoxide reduction of permanganate is seen to be analo-gous in some respects to the formic acid reduction. In both cases the postu-lated mechanisms involve attack on the carbon atom by a permanganate oxygen. The reactions of carbon monoxide with Ag(l) and Hg(ll) seem closer to those of hydrogen than to those of formate. This is probably related to the nature of the intermediates? those of hydrogen and carbon monoxide being more covalent than those of formate ion. The catalytic reduction of permanganate by carbon monoxide is not like the catalytic reduction by hydrogen. Rather i t is very close to the carbon monoxide reductions of Hg(ll) and Ag(l). The key to these reductions seems to be the formation of carbon monoxide insertion complexes. 128. Both Ag(l) and Hg(Il) are d 1 0 Ions as are Cu(l) which forms quite stable CO complexes and Au(l) which reportedly oxidizes CO (105). Orgel (147) has pointed out that the d - s separation of these ions i s low enough to permit extensive d - s mixing. He a t t r i b u t e s the f a c t that these ions tend to form l i n e a r complexes to t h i s mixing. For the i s o e l e c t r o n i c ions Z n ( l l ) and C d ( I l ) the energy differences are too large f o r t h i s to take place. The T l ( l l l ) ion i s intermediate having a d - s separation much greater than that of H g ( l l ) , but somewhat l e s s than that of Z n ( l l ) . Neither C d ( l l ) nor T l ( I I l ) have any e f f e c t on CO and i t i s u n l i k e l y that Z n ( I l ) i s active. The ions which are active towards CO presumably donate d-electrons to form p i bonds and can then form reasonably strong metal-carbon sigma bonds through the d - s h y b r i d i z a t i o n . Comparison of the I s o e l e c t r o n i c Species CO, N 2 and CN'' Carbon monoxide resembles nitrogen i n i t s physical properties, but i s c l o s e r to cyanide i n chemical behaviour. The r e l a t i v e chemical a c t i v i t y of the three species can be i l l u s t r a t e d by reference to the ease of the ad-d i t i o n of an oxygen atom. N 2 + &) 2 > N 20 A F ° = 24.8 kcal/mole CO + ^P2 > C0 2 A F ° = -61.4 kcal/mole CN" + i o 2 > OCN" A F ° = -63.2 kcal/mole J a f f e and Orchin (95) have given an explanation of the d i f f e r e n c e s i n the r e a c t i v i t i e s of nitrogen and carbon monoxide. The 2s and 2 p x (x i s the interatomic axis) atomic o r b i t a l s of the nitrogen atom can hybridize or mix to form two sigma o r b i t a l s one with a small amount and one with a large amount of p- character. These new o r b i t a l s 129. will have energies somewhat above the original s- level and below the original p- level respectively. The energies of the remaining pi orbitals are unchanged from the original p- levels. The hybridized atomic orbitals of nitrogen will contain two electrons in the lowest sigma orbital, one in the next sigma orbital and two in the doubly degenerate pi orbitals. When two nitrogen atoms combine the f i r s t sigma orbitals remain unchanged and may be identified with the lone pairs of electrons on the two nitrogen atoms. The second sigmse orbitals combine to form a strong sigma bonding molecular orbital and a high energy, vacant, sigma antibonding molecular orbital.. Similarly the pi orbitals combine to form a degenerate pair of pi bonding orbitals and a degenerate pair of pi antibonding orbitals. The relative levels of these in increasing order are bonding sigma, non-bonding sigma, bonding pi al l fully occupied. The lowest unoccuDied orbital is an anti-bonding pi orbital. The points of interest are that the lone pairs being largely of s- character have no directional tendencies. The highest occupied orbital is a bonding orbital so that electron promotion or ionization will weaken the interatomic bond. The lone pairs being in low lying levels will have l i t t l e or no tendency to take part in sigma bonding. In the formation of the carbon monoxide molecule the atomic orbitals of carbon and oxygen are also hybridized. Because of the differ*-ences in atomic charges the situation will not be the same as with nitrogen. The lowest oxygen sigma orbital will be considerably below the lowest carbon sigma orbital so these will not combine. The lowest carbon sigma orbital will contain one electron and will combine with the second oxygen sigma orbital also containing one electron to form the sigma bonding and anti-bonding orbitals. The pi orbitals will combine as before to form bonding 130. and antibonding o r b i t a l s , but the bonding p i o r b i t a l w i l l l i e below the second sigma o r b i t a l of carbon. This l a t t e r o r b i t a l w i l l contain the carbon lone p a i r and being l a r g e l y p i n character w i l l be strongly d i r e c t i o n a l . Also with carbon monoxide the highest occupied o r b i t a l being non-bonding, i o n i z a t i o n w i l l not weaken the carbon - oxygen bond. The cyanide ion w i l l be q u a l i t a t i v e l y s i m i l a r to carbon monoxide. The main differences are that as carbon and nitrogen are c l o s e r than carbon and oxygen the separation of the l e v e l s w i l l not be as pronounced. Further-more, the cyanide ion with a lower nuclear charge w i l l have one electron held much l e s s strongly than i n carbon monoxide. The i o n i z a t i o n p o t e n t i a l s of the species are* Ng 359 k c a l ; CO 324. k c a l ; CN* 103 k c a l . These f i g u r e s are i n accord with the r e l a t i v e reac-t i v i t i e s . Another f a c t o r i n considering the r e a c t i v i t i e s i s the energy d i f -ferences between the non-bonding sigma o r b i t a l s and the antibonding p i o r b i t a l s . Both carbon monoxide and cyanide complexes are s t a b i l i z e d to varying degrees by back bonding to t h e i r vacant p i o r b i t a l s . E l ectron donation from a l i g a n d sigma o r b i t a l and back donation to a p i o r b i t a l i s equivalent to an e l e c t r o n t r a n s i t i o n between these l e v e l s . The pertinent values are* N 2 168 k c a l ; CO 138 k c a l ; CN" 115 k c a l . These again are i n the order of t h e i r r e a c t i v i t i e s . 131. BIBLIOGRAPHY 1. J . Halpern, J . Phys. Chem. 63, 398 (1959). 2. B. J . Zwolinski, R. J . Marcus and H. Eyring, Chem. Rev., 157 (1955). 3. F. Basolo and R. G. Pearson, "Mechanisms of Inorganic Reactions, John Wiley and Sons, Inc., New York, 1958, Chapt. 7. 4. H. Taube, i n "Advances i n Inorganic Chemistry and Radiochemistry", Vol. 1, Ed. H. J . Emelius and A. G. Sharpe, Academic Press, Inc., New York, 1959, p. 1. 5. P. George and J . S. G r i f f i t h , i n "The Enzymes", Vol. 1, 2nd ed., Ed. Boyer, Lardy and R. Myrback, Academic Press Inc., New York, 1959, Chapt. 8. 6. D. R. Stranks, i n "Modern Coordination Chemistry", Ed. J . Lewis and R. G. Wilkins, Interscience Publishers Inc., New York, I960, Chapt. 2. 7. J . Halpern, Quart. Rev., 15_, 207 (1961). 8. J . C. Sheppard and A. C. Wahl, J . Am. Chem. Soc. 72, 1020 (1957). 9. H. Taube, H. Myers and R. L. Rich, i b i d . , 75_, 4118 (1953). 10. H. Taube, H. Myers, i b i d . , 76, 2103 (1954). 11. A. Haim and W. K. Wilmarth, i b i d . , 82, 509 (1961). 12. J , Silverman and R. W. Dodson, J . Phys. Chem., j>6, 846 (1952). 13. R. W. Dodson and N. Davidson, i b i d . , j>6, 866 (1952). 14. J . Hudis and R. W. Dodson, J . Am. Chem. Soc. 78, 911 (1956). 15. A. E. Ogard and H. Taube, i b i d . , 80, 1084 (1958). 16. R. K. Murmann, H. Taube and F. A. Posey,ibid., 72, 262 (1957). 16A., M. Ardon, J . Levitan and H. Taube, i b i d . , 84., 872 (1962). 17. A. Zwickel and H. Taube, i b i d . , 83, 793 (1961). 18. H. Taube, i b i d . , 77, 4481 (1955). 19. A. M. Zwickel and H. Taube, Disc. Faraday Soc., 22 , 42 (i960). 20. N. A. Bonner and J . P. Hunt, J . Am. Chem. S o c , 82, 3826 (i960). 21. W. F. Libby, J . Phys. Chem. ^6, 863 (1952). 22. R. J . Marcus, B. J . Zwolinski and H. Eyring, i b i d . , j>8, 432 (1954). 132. 23, D. Cohen, J. C. Sullivan and J. C. Hindman, J. Am. Chem. Soc, 76, 352 (1954). 24, D. Cohen, J. C. Sullivan, E. S. Amis and J. C. Hindman, ibid., 78, 1543 (1956). 25. R. A. Marcus, J. Chem. Phys., 2£, 966 (1956). 26. R. A. Marcus, ibid., 26, 867 (1957). 27. R. A. Marcus, ibid., 26, 872 (1957). 28. R. A. Marcus, Disc. Faraday Soc., 29_, 21 (i960). 29. N. s . Hush, Trans. Faraday Soc, £7, 557 (1961). 30, R. E. Kirk and A. W. .Browne, J. Am. Chem. Soc. j>0, 337 (1928). 31. V. C. E. Higginson, D. Sutton and P. Wright, J. Chem. Soc. 1380 (1953). 32. w. C. E. Higginson and D. Sutton, ibid., 1402, (1953). 33. w. C. E. Higginson and P. Wright, ibid., 1551 (1955). 34. w. C. E. Higginson and J. Marshall, ibid., 447 (1957). 35. L. Michaelis, Cold Spring Harbor Symposia Quant. Biol., 7, 33 (1939). 36. F. H. Westheimer, in "The Mechanism of Enzyme Action" Ed. W. D. McElroy and B. Glass, Johns Hopkins Press, Baltimore, 1954, p. 321, 37. F. Basolo and R. G. Pearson, Reference 3, p. 305. 38. H. Taube, Reference 4, p. 3. 39. P. A. Shaffer, J. Am. Chem. Soc. 5J>, 2169 (1933). 40. P. A. Shaffer, Cold Spring Harbor Symposia Quant. Biol., 7, 50 (1939). 41. A. E. Remick, J. Am. Chem. Soc, 62, 94 (1947). 42. J. Halpern, Can. J. Chem., 37, 148 (1959). 43. J. Halpern, E. R. MacGregor and E. Peters, J. Phys, Chem., 60, 1455 (1956). 44. J. Halperin and H. Taube, J. Am, Chem. Soc, 74_, 375 (1952). 45. R. Stewart, ibid., 72, 3057 (1957). 46. B. J. Masters and L. L. Schwartz, ibid., 83, 2620 (1961). 47. E. Rona, ibid., 72, 4339 (1950). 48. T. W. Newton, J. Phys. Chem., 62, 1493 (1959). 49. T. w. Newton, ibid., 6j2, 943 (1958). 133. 50. J . C. S u l l i v a n , A. J . Zi e l e n and J . C. Hindman, J . Am, Chem, S o c , 82, 5288 (1960). 51. R. H. Betts, Can. J . Chem., 33_, 1780 (1955). 52. F. B. Baker, T. W. Newton and M. Kahn, J . Phys. Chem., 6£, 109 (i960). 53. J . Halpern and J . G. Smith, Can. J . Chem., 34., 1419 (1956). 54. F. B. Baker and T. W. Newton, J . Phys. Chem., 6j>, 1897 (1961). 55. E. A. Kanevskii, L. A. Fedorova, Zh. Neorg. Khim., 2216 (i960). 56. J , J . Katz and G. T. Seaborg, "The Chemistry of the Actinide Elements", Methuen and Co., Ltd., London, 1957. 57. T. W. Newton and F. B. Baker, Inorg. Chem., 1, 368 (1962). 58. G, Gordon and H. Taube, i b i d . , 1, 69 (1962). 59. E. Roig and R. W. Dodson, J . Phys. Chem. 65_, 2175 (1961). 60. S. W. G i l k s and G. M. Waind, D i s c Faraday S o c , 22, 102 (i960). 61. B. Baysal, Actes Intern. Congr. Catalyse, 2 e, Paris I960, 1, 559 (Pub. 1961). 62. L. G. Carpenter, M. H. Ford-Smith, R. P. B e l l and R. W. Dodson, D i s c Faraday Soc., 2<?, 92 (i960).. 63. C. H. Brubaker, J r . , and J . P. Mickel, J . Inorg. Nuclear Chem., ^, 55 (1957). 64. C. H. Brubaker, J r . , K. 0. Groves, J . P. Mickel and C. P. Knop, J . Am. Chem. S o c , 22, 4641 (1957). 65. C E. Johnson, J r . , i b i d . , 74., 959 (1952). 66. K. G. Ashurst and W. C. E. Higginson, J . Chem. Soc., 3044 (1953). 67. F. R. Duke and B. Bornong, J . Phys. Chem., 60, 1015 (1956). 68. D. H. Ir v i n e , J . Chem. S o c , 1841 (1957). 69. K. G. Ashurst and W. C. E. Higginson, J . Chem. S o c , 343 (1956). 70. A. M. Armstrong and J . Halpern, Can. J . Chem., 3J5, 1020 (1957). 71. M. Ardon and R. JL Plane, J . Am. Chem. Soc., 81, 3197 (1959). 72. W. C. E. Higginson, D. R. Rosseinsky, J . B. Stead and A. G. Sykes, Disc. Faraday Soc., 2_2, 49 (i960). 73. A. G. Sykes, J . Chem. Soc., 5549 (1961). 74. J . W. Gryder and M. C, Dorfman, J . Am. Chem. S o c , 83, 1254 (1961).. 134. 75. A. M. Armstrong and J . Halpern, unpublished observations. 76. H. N. Halvorson and J . Halpern, J . Am. Chem. S o c , 78 , 5562 (1956). 77. J . Halpern and S. M. Taylor, Disc.Faraday Soc., 29_, 174 (i960). 78. C. H. Brubaker, J r . , i n "Advances i n the Chemistry of Coordination Compounds". Ed., S. Kirschener, MacMillan, New York, 1961, p. 117. 79. A. C. Harkness and J . Halpern, J . Am. Chem. Soc., 81, 3526 (1959). 80. A. A. Fr o s t and R. G. Pearson, "Kinetics and Mechanism" John Wiley & Sons, Inc., New York, 1953, p. 17. 81. G. Biederman, Arkiv Kemi £, 441 (1953). 82. K. A. Kraus and F. Nelson, J . Am. Chem. S o c , 72 , 390 (1950). 83. K. A. Kraus and F. Nelson, i b i d . , 77, 3721 (1955). 84. R. H. Betts, Can. J . Chem., 22, 1775 (1955). 85. " S t a b i l i t y Constants, Part I I " , Chem. Soc. Special Publ., No. 7 (London, 1958). 86. W. M. Latimer, "The Oxidation States of the Elements and Their Potentials i n Aqueous Solutions", 2nd ed., P r e n t i c e - H a l l , Inc., New York, 1952. 87. T. W. Newton and S. W. Rabideau, J . Phys. Chem., 62, 365 (1959). 88. J . Halpern, Quart. Rev., 10, 463 (1956). 89. J . Halpern, Advances i n C a t a l y s i s 11, 301 (1959). 90. Y. K. Syrkin and M. E. Dyatkina, "Structure of Molecules and the Chemical Bond", Butterworths, London 1950, p. 137. 91. I . Langmuir, J . Am. Chem, S o c , 4JL, 1543 (1919). 92. R, S. Mulliken, Rev. Mod. Phys., 1 (1932). 93. W. M o f f i t t , Proc. Roy. Soc., A196, 524 (1949). 94. R. C. Sahni, Trans. Faraday S o c , 42, 1246 (1953). 95. H. H. J a f f e and M. Orchin, Tetrahedron, 10, 212 (I960). 96. C. Mond, C. Langer and F. Quincke, J . Chem. S o c , j>7, 749 (1890). 97. J . W. Richardson, i n "Organometallic Chemistry", Ed. H. Zeiss, Reinhold, New York, I960. Chapt. 1. 98. J . Chatt, P. L. Pauson, L. M. Venanzi, i b i d . , Chapt. 10. 99. I. Wender, H. W. Sternberg and M. Orchin, i n " C a t a l y s i s " Vol. 5, Ed. P. H. Emmett, Reinhold, New York, 1957, p. 73. 135. 100. H. W. Sternberg and I. Wender, International Conference on Coordination Chemistry, Chem. Soc. Special Publ. No. 13, London, 1959, p. 35. 101. T. H. Coffield, R. D. Closson, J. Kozikowski, J. Org. Chem., 22, 598 (1957) . 102. F. Calderazzo and F. A. Cotton, Inorg. Chem., 1, 30 (1962). 103. M. Katz, Advances i n Catalysis j>, 177 (1953). 104. J. R. Dixon and J. E. Longfield in "Catalysis" Vol. 7, Ed. P. H. Emmett, Reinhold, New York, I960. 105. J. W. Mellor, "A Comprehensive Treatise on Inorganic and Theoretical Chemistry", Longmans Green & Co., London, 1923, pp 909-43. 106. J. Donan, Monatsh 26, 525 (1905). 107. J. Donan, ibi d . , 27, 71 (1906). 108. F. C. Phi l l i p s , Am. Chem. J. 16, 163 (1894). 109. F. C. Ph i l l i p s , i b i d . , 16, 255 (1894). 110. C. Winkler, Z. Anal. Chem., 28, 269 (1889). 111. L. M. Dennis and C. G. Edgar, J. Am. Chem. Soc. 12, 859 (1897). 112. F. Jean, Compt. rend, 135, 746 (1902). 113. G. Just and Y. Kauko, Z. physik Chem., 82, 71 (1913). 114. K. A. Hofmann, German Patent 307, 614 (1919). 115. A. Mermet, Compt. rend. 124, 621 (1897). 116. G. Bauch, F. Pawlek and K. Plieth, Z. Erz. und Metall., 11, No. 11, 1 (1958) . 117. A. H. Webster, "A Kinetic Study of Some Homogeneous Reactions of Mole-cular Hydrogen with Metal Ions in Aqueous Solution", PhD Thesis, Department of Metallurgy, University of British Columbia, 1957. 118. Seidel, "Solubilities of Inorganic and Metal Organic Compounds", 3rd ed., Vol. 1, 1940, p. 217. 119. J. Halpern, J. F. Harrod and P. E. Potter, Can. J. Chem., 37, 1446 (1959). 120. L. G. S i l l e n , Quart. Rev., 13, 146 (1959). 121. K. B. Wiberg and R. Stewart, J. Am. Chem. Soc, 78, 1214 (1956). 122. G. Friedlander and J. W. Kennedy, "Nuclear and Radiochemistry", John Wiley ft Sons, Inc., New York, 1955. 123. G. J. Korineck and J. Halpern, J. Phys. Chem., 60, 285 (1956). 136. 124. W. Schoeller, W. Schrauth and W. Essens, Ber., 46, 2864 (1913). 125. W. Manchot, ibid., £3, 984 (1920). 126. J. Halpern and S. F. A. Kettle, Chem. and Ind., 668 (1961). 127. E. Lange and H. Sattler, Z. physik. Chem., A179, 427 (1937). 128. J. Halpern and A. C. Harkness, J..Chem. Phys. 31, 1147 (1959). 129. J. Bigeleisen, ibid., 32.> 1583 (I960). 130. A. L. W. Kemp, "Kinetics of the Decomposition of Methoxycarbonylmercuric Chloride by Hydrochloric Acid", B.Sc. Thesis, Department of Chemistry, University of British Columbia, 1961. 131. A. C. Harkness and J. Halpern, J. Am. Chem. Soc, 83_, 1258 (1961). 132. R. T. McAndrew, "Carbon Monoxide Reduction of Aqueous Silver Acetate", PhD Thesis, Department of Metallurgy, University of British Columbia, 1962. 133. S. Nakamura and J. Halpern, J. Am. Chem. Soc, 83, 4102 (1961). 134. S. Nakamura, "Reduction of Silver Amine Complexes", MSc Thesis, Depart-ment of Chemistry, University of British Columbia, 1962. 135. J. J . Byerley and E. Peters, Paper presented at the Annual Meeting of the A.I.M.M.E., Dallas, Feb., 1963. 136. A. Carrington and M. C. R. Symons, J. Chem. Soc., 3373 (1956). 137. H. Lux, ZT, Naturforsch, 1, 281 (1946). 138. T. Freund, J. Inorg. Nuclear Chem., 15_, 371 (i960). 139. R. Stewart and R. Van der Linden, Can. J. Chem., 38, 2237 (I960). 140. E. Peters and J. Halpern, J. Phys. Chem., jg, 793 (1955). 141. A. H. Webster and J. Halpern, ibid., 61, 1239 (1957). 142. A. H. Webster and J. Halpern, Trans. Faraday Soc, j>3, 51 (1957). 143. C. E. H. Bawn and A. 0. White, J. Chem. Soc, 339 (1951). 144. s. M. Taylor and J. Halpern, J. Am. Chem. Soc, 81, 2933 (1959). 145. J. Halpern, Advances in Catalysis 11, 301 (1959), 146. A. Carrington, D. Schonland and M. C. R. Symons, J. Chem. Soc 659 (1957). 147. L. E, Orgel, J. Chem. Soc., 4186 (1958). 

Cite

Citation Scheme:

        

Citations by CSL (citeproc-js)

Usage Statistics

Share

Embed

Customize your widget with the following options, then copy and paste the code below into the HTML of your page to embed this item in your website.
                        
                            <div id="ubcOpenCollectionsWidgetDisplay">
                            <script id="ubcOpenCollectionsWidget"
                            src="{[{embed.src}]}"
                            data-item="{[{embed.item}]}"
                            data-collection="{[{embed.collection}]}"
                            data-metadata="{[{embed.showMetadata}]}"
                            data-width="{[{embed.width}]}"
                            async >
                            </script>
                            </div>
                        
                    
IIIF logo Our image viewer uses the IIIF 2.0 standard. To load this item in other compatible viewers, use this url:
http://iiif.library.ubc.ca/presentation/dsp.831.1-0062052/manifest

Comment

Related Items