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The kinetics of carbon monoxide absorption in basic solutions at elevated temperature McDonald, Robert Douglas 1964

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THE KINETICS OF CARBON MONOXIDE ABSORPTION.IN BASIC- SOLUTIONS AT ELEVATED. TEMPERATURE by • ROBERT DOUGLAS McDONALD " A THESIS: SUBMITTED'IN: PARTIAL FULFILMENT OF THE REQUIREMENTS FOR. THE DEGREE.OF MASTER OF SCIENCE IN THE DEPARTMENT of METALLURGY We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA October 1963 . In presenting . t h i s .thesis i n p a r t i a l f u l f i l m e n t of the requirements f o r an advanced degree at the U n i v e r s i t y of B r i t i s h Columbia, I agree th a t the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r reference and study. I f u r t h e r agree that permission f o r extensive.copying of t h i s t h e s i s f o r s c h o l a r l y purposes may be granted by the Head of my Department or by h i s r e p r e s e n t a t i v e s . I t i s understood tha t copying or p u b l i c a t i o n of t h i s t h e s i s f o r f i n a n c i a l g a i n s h a l l not be allowed without my w r i t t e n permission. Ifetallurgy Department of The U n i v e r s i t y of B r i t i s h Columbia, Vancouver 8, Canada. October 25, 1963 Date • ABSTRACT . The.kinetics of the absorption of carbon monoxide, by basic solutions was studied at 80°C and. carbon-monoxide pressures up. to 30 atmos-pheres. •The reaction was followed by the rate of decrease of carbon monoxide.pressure i n a closed system. The observed k i n e t i c s in. potassium hydroxide solutions y i e l d a rate law of the form d[C04. 1 tot dt .. - l i g Q = d[CH0 2-] = k r.[C0 s][0H"] dt . a i s s + + + No influence from: L i } - Na } K ions was detected and no c a t a l y t i c e f f e c t from Ag(I),- Cu(II), T1(I), N03~, Mn04~ was observed. The k i n e t i c s are consistent with a mechanism which includes the i n s e r t i o n of a carbon monoxide molecule i n t o the hydroxyl bond,,viz. . 0' C 0 d i s s + OH' ^C = 0 i - > 0'-f a s t H - C 0 0 — The r a t e - c o n t r o l l i n g step above 90°C was found ; to be the mass t r a n s f e r of carbon monoxide from the gas phase i n t o the l i q u i d phase under the conditions involved i n . t h i s study. i i i . ACKNOWLEDGEMENT The author wishes to express h i s sincere thanks to Dr. E. Peters for h i s i n s p i r i n g guidance and encouragement throughout the period of study. He also g r a t e f u l l y acknowledges the f i n a n c i a l assistance given to him, i n the form of a Bursary, by the National. Research Council of Canada. i v . TABLE, OF CONTENTS Page INTRODUCTION 1 • EXPERIMENTAL 5 Apparatus 5 - Materials 9 Procedure 9 Analysis of Results 9 RESULTS AND DISCUSSION Ik -A. Reaction Controlled Region . ik . . I d e n t i f i c a t i o n of Products ik Stoichiometry of Reaction 16 K i n e t i c s 17 E f f e c t of Formate 21 E f f e c t of Cation 23 E f f e c t of Ionic Strength 23 E f f e c t of Catalysts 2k Rate i n Buffer Solution . . ' 26 E f f e c t of Temperature 32 The Reaction Mechanism . . 35 •B. . D i f f u s i o n C o n t r o l l e d Region kO K i n e t i c s . . . . . . .kO •Effect of Temperature k2 C. Ap p l i c a t i o n of Formate Reaction . kk . CONCLUSION . k6 APPENDIXES . . . . . 7^ A. Observations Concerning Reduction of Cobalt (II)by Carbon Monoxide 7^ B. Summary of Experimental Data 50 REFERENCES 51 V. LIST OF FIGURES Page Figure 1. Energy .Levels i n N 2 and"in-CO. 2 Figure 2. Titanium:Autoclave, 6 Figure 3- Photograph of- Shaking Autoclave. 7 Figure k. Schematic. IJ'rawing of Reactor System. . . . . . . . . . . . . . . 8 Figure 5. -Moles of Contained CO as a Function of CO Pressure 10 • Figure 6. Typical.Pressure-Time Curves. . . 12 Figure 7. T i t r a t i o n Curve of F i n a l Formate Solution of Run Number 9. 15 Figure 8. V a r i a t i o n of Contained Gas'Constant C with Temperature. ... ..19 Figure 9« T y p i c a l Second Order'Rate Plots 22 Figure 10. F i r s t Order Rate :Plot i n Phosphate Buffer. .. 28 •Figure 11. Heat o f Ionization of Mono-hydrogen Phosphate. .31 • Figure 12. - Arrhenius Plot of I n i t i a l Reaction Rate . . . . 33 Figure 13. F i r s t Order Rate Plots i n Mass Transfer Controlled Region.. hi Figure Ik. Second.Order Rate Plots i n Mass Transfer-Controlled Region. hj> v i . LIST OF. TABLES Page I. T y p i c a l Analysis of Pressure-Time Curve 13 I I . V e r i f i c a t i o n of Stoichiometry of Run Number 9 16 I I I . V a r i a t i o n of S o l u b i l i t y of Carbon Monoxide, and Contained Gas , Constant with Temperature 20 IV. Experimental Rate Constants at 80°C. . . . . . . . . . . . . . . 20 V. ' E f f e c t of Change of A l R a i i Metal Cation. . 23 VI. E f f e c t of Ionic Strength . ' 2k VII. E f f e c t of Ions as Catalysts on Rate of Formation of Formate S a l t s . 2k VIII. • Arrhenius Plot Data ~$k -IX. Vibrational. Stretching Frequencies 38 X. V a r i a t i o n of D i f f u s i o n C o e f f i c i e n t s and I n i t i a l D i f f u s i o n Rates with Temperature.. . ... . . . . ... . ...... k2 INTRODUCTION Carbon monoxide has been used to p r e c i p i t a t e metals from aqueous solutions. Several d e t a i l e d studies have-been c a r r i e d out to determine 1-5 the mechanisms of some of these reactions . Much of the work, p a r t i c u l a r l y the high, temperature, high pressure reduction of s i l v e r and copper, has been studied i n a c i d i c solutions where absorption of carbon monoxide due to formation of formate s a l t s was found to be n e g l i g i b l e . -In a l l the carbon monoxide reduction studies involving metal ions i t was postulated that the mechanism involved intermediates that could be described as i n s e r t i o n complexes, i . e . 0 II CO + -MOR — 3 — -M-C-OR . . . . . ( 1 ) where M i s a metal ion and -OR i s a basic ligand (e.g..OH., H 20, OAc ). The CO molecule breaks the metal-to-ligand bond and becomes "inserted" 6 between the metal and ligand. •  Schoeller i s o l a t e d a.stable' complex of t h i s type from a methanolic mercuric acetate s o l u t i o n , the structure of 7 which was determined to be 0 I AcO-Hg-C-OCH3 Carbon monoxide i s i s o e l e c t r o n i c with nitrogen gas. • Langmuir postulated the structure _ + :C ^ = 0: f o r carbon monoxide i n analogy to nitrogen - 2 -As i s w e l l known, carbon monoxide has a very small dipole moment and oxygen i s more electrophi l i c.; than carbon, both of which are inconsistent wi th the above s tructure. In terms- of Paulings resonant structures three configurations can be drawn. + r C-0 c=o + 0 In the more exact notation of molecular o r b i t a l theory the ground state configuration for both CO and N 2 i s given as: K . K . . ( t o - ) (S < 7 ~ r ( u c r ) . ( TT x ) . The reason for the inertness- of nitrogen and. the r e l a t i v e l y high r e a c t i v i t y of carbon monoxide, p a r t i c u l a r l y as a nucleophllie•: reagent, can 8. be understood by reference to the energy l e v e l diagram of Jaffe and Orchin . N N / XXX i U xx, 1/ V» / I — ( / \ \XXXX/ , \. \ rr \ crsc- . / \ \ \ h xx xx. CO •I \ I \ ' \ r—w \, , * w >. * x x ____xxxx XX j- v / \ / \ ' M \. I \ \ i i. xx xx ,' \ Figure 1. Energy Levels i n N 2 and i n CO. In nitrogen the outermost electrons are bonding.electrons and r e l a t i v e l y unreactive. In carbon monoxide, the outer electrons are non-bonding and therefore more reactive,, p a r t i c u l a r l y toward e l e c t r o -9 phrles. As c a l c u l a t e d by Sahni these electrons are l o c a l i z e d as a lone p a i r on the carbon end* of the molecule. This.explains the r e a c t i v i t y of the carbon atom: in.the molecule. The non-bonding ;pair of electrons l o c a l i z e d on the oxygen atom are the low energy electrons and very i n e r t . The small dipole moment of CO e x i s t s because the e l e c t r o n density about 9 the carbon atom, i s not centered on the carbon nucleus J but s l i g h t l y outside . Carbon monoxide, although a stable gas, is. thermodynamically unstable r e l a t i v e to carbon and carbon dioxide. 2C0 C + C0 2 The k i n e t i c s of t h i s r e a c t i o n at ordinary temperatures are so slow that the reaction can be ignored. Carbon monoxide a l s o undergoes the water gas r e a c t i o n : CO + H 20 —*-.C0 2 + H 2 .(2) 10 Royen has postulated that t h i s reaction proceeds by way of an intermediate formate ion. Although carbon•monoxide i s the anhydride of formic a c i d , i t "does not react with water to give formic a c i d except at unusually high pressures. Reference to the pH-potential diagram of the CO, CH0 2", G0 2 system"^", shows that CO i s unstable i n base r e l a t i v e to formate while i n a c i d the opposite e f f e c t i s true. - k -The re a c t i o n of CO with s o l i d a l k a l i to give" formate s a l t s has 12 been known since the mid-19th Century „ The addition of water increased the rate of reacti o n . •Above 220°C carbonate and hydrogen i s formed rather than formate. The synthesis of formic a c i d or formate s a l t s from CO and aqueous a l k a l i s has been.the subject of several papers and • ' • 13 -15 patents . The synthesis of cesium formate from fused cesium hydroxide 15 has.been reported . .The temperature i s f a i r l y high- (280-300°C ). I f the temperature i s above 300°C cesium ©xalate i s formed. Sodium formate a l s o forms from the reaction between C0 2 and sodium, amalgam . Analogous reactions using alcohols have been reported^ methyl alcohol and CO y i e l d s 17 18 methyl acetate or a c e t i c a c i d , i f catalysed with BF 3 . Formamide •7° ' 19 HC -NH 2 has been synthesized from ammonia and carbon monoxide i n glycerine . 20 ' , S i r o t k i n has analyzed the reaction thermodynamically and shown that r e a c t i o n w i l l . o c c u r to zero CO pressure. .Kodama et a l . have studied 21 22 the reaction of CO with NaOH and N H 4 O H and report a c t i v a t i o n energies • of 22 Kcal/mole and I7.8 Kcal/mole r e s p e c t i v e l y . The dehydration of formic a c i d r e l e a s i n g carbon monoxide i s w e l l 23 2k known. I t can be achieved with strong H2S04 and with s i l i c a g e l .The object of the present work was to resolve the k i n e t i c s and mechanism of the carbon monoxide absorption by basic, solutions that lead to the formation of formate s a l t s . No complete study has been done previously. The complete r e s o l u t i o n of t h i s reaction w i l l a i d i n the reduction k i n e t i c s of metal ions from basic s o l u t i o n , where i t frequently i s an important side reac t i o n . - 5 -- EXPERIMENTAL Apparatus A titanium shaking autoclave of about 103 cc capacity, s i m i l a r to that described by McAndrew , was used in. t h i s study (Figures 2 and 3). A g i t a t i o n was achieved by shaking.the reactor h o r i z o n t a l l y at 275 o s c i l -l a t i o n s per minute through a l.'l/2 inch stroke. Pressure measurement was by use of a Consolidated Electrodynamics Corporation pressure transducer pickup, (Type 4-311) mounted i n the closed system. • Heating was by a band heater wrapped around the autoclave. The heater, rated at kOO watts and 110 v o l t s , was connected by a shunted on-off re l a y of a temperature c o n t r o l l e r to a variac Transformer. The temperature was c o n t r o l l e d by a thermistor sensing element mounted i n a thermal well i n the center of the autoclave and a.Yellow Springs Instruments Thermistemp C o n t r o l l e r , Model 71. With the r e l a y shunted with a k ohm r e s i s t o r , to give a lower heating rate i n the o f f p o s i t i o n , and appropriate adjustment of the V a r i a c , the temperature c o n t r o l was ± 0.15°.'C. A.schematic drawing of the system i s shown i n Figure k. A Beckman model G pH meter with a model 391^2 combination g l a s s -calomel electrode was used f o r a l l pH t i t r a t i o n s . • A s p e c i a l p i p e t t e was constructed to charge a constant l i q u i d volume to the autoclave. The.volume of t h i s pipette was 83.h ml at room temperature. -6 -1/8" Ermeto Set Screw (8 at ) Cover (titanium) Thrust Washer Cover Nut - A5136 S t e e l Titanium to Titanium Seal Gasket/ Teflon Thermowell •Autoclave Body, Titanium 1/8" Ermeto l / l 6 " Ermeto (k at 90°) Figure 2. . Titanium.Autoclave ( f u l l scale) - 7 -Figure 3. Photograph of Shaking Autoclave - 8 -R A Q-4 G 4 A. -Autoclave C . - Temperature C o n t r o l l e r G. CO Gas Cylinder R. Recorder T. .Transducer V-|_ Gas - I n l e t Valve Vg - Gas Outlet Valve Figure k. • Schematic Drawing of Reactor System - 9 -Materials .Chemicals were of reagent grade and were used without f u r t h e r p u r i f i c a t i o n . Analysis of potassium hydroxide and phosphate bu f f e r solutions was by pH t i t r a t i o n using standard HC1. Carbon monoxide .(99-5$) was supplied by Matheson Company. Chromatographic analysis indicated that N 2 , O 2 and H 2 were present i n quantities l e s s than 0.2$, 0.01$ and 0.002$ r e s p e c t i v e l y . Procedure A constant volume of s o l u t i o n (83.k cc from s p e c i a l pipette) was charged to the autoclave which was then sealed. A small amount of r e s i d u a l a i r was- l e f t i n the system*. The autoclave was heated to temperature and CO gas to the desired pressure admitted from two sides, thus f l u s h i n g o u t l e t and i n l e t l i n e s i n t o autoclave. The system was resealed and the pressure drop, recorded on an• E. H. Sargent and Company Model SR,-Chart Recorder, (see Figure 6.) Analysis of Results C a l i b r a t i o n curves of moles of contained gas as a function of CO pressure are shown i n Figure 5- These were obtained experimentally by measuring, at room, temperature and atmospheric pressure, the volume of gas contained i n the autoclave, at various pressures and temperatures, and d i v i d i n g by the known molar volume f o r ..CO under these room conditions. A 0 2 was added at the end of one run. A slow pressure drop indicated a . s l i g h t r e a c t i o n with formate but the rate was much slower than the formate s a l t formation. -11 -These c a l i b r a t i o n s are str a i g h t l i n e s passing, through the o r i g i n and consistent with an equation of state of the form p (V + c) = nRT The constant c i s not zero because the van der Waals volume of the gas phase, the equivalent gas volume of the l i q u i d phase, and a small, volume i n the l i n e to the transducer, were not taken i n t o account. The extent of reaction can be calculated from.the pressure drop curves (Figure 6) and these c a l i b r a t i o n s . A t y p i c a l a n alysis i s given i n Table I. .Rate constants were obtained from the slope of second order p l o t s containing-the function . ( N C Q ) ± - [0H _] - + , [ O H " ] .V I Join p l o t t e d against time. TABLE.1. . Typical. Analysis of Pressure-Time Curve -Experiment Number.7,, [ KOH]^ = 0.168 M/ T = 80°C, V± = 83.4 cc. Time p c o P t o t - P 0 rC0 NC0 =  PC0 x c Extent of Reaction Moles.OH" of s o l u t i o n [OH"] 0 0 1 - [OH ]±+[QR-] ( N c o ) i-[OH-].+[OH"] In 1 (min) (mv) (atm) (moles) (moles CO) (M) roH"i [OH"] 0 12.62 19.9^ 0.0184 0 0.0140 0.168 1.32 0.274 .5 11.11 17.55 0.0162 0.0022 0.0118 0.142 1.37 0.317 10 10,15 16. Ok 0.0148 0.0036 0.0104 0.125 1,42 0.354 20 8.72 13.78 0.0127 O.OO57 O.OO83 0.100 1.53 0.425 30 7.80 12.32 0.0114 0.0069 0.0071 O.O85 1.62- 0.484 4o .7.07 11.17 0.0103 0.0081 O.OO59 0.071 1.75 0.558 50 6.U9 10.25 O.OO95 0.0089" 0.0051 0.061 1.87 0.624 60 6.03 9-53 0.0088 0.0096 0.0044 O.O53 2.00 0.693 70 5.65 8.93 0.0082 0.0102 0.0039 0.046 2,15 O.766 80 5.32 •841 O.OO78 0.0106 0.0034 0.040 2.33 0.844 90 5.O5 7.96 0.0074 0.0111 0.0030 0.035 2.51 0.921 . 100 4.80 7.58 0.0070 0.0114 0.0026 0.031 2.71 0.999 -Ik -RESULTS AND DISCUSSION The reaction of carbon monoxide with hydroxyl ions i s considered as a two step reaction C 0 g C 0 a q _ . . . . - .(3a) CO + OH" — > - C H 0 2 (3b) Carbon monoxide i n the gas phase must dissolve i n the aqueous s o l u t i o n . The dissolved CO must then d i f f u s e to, and react with the hydroxyl ions i n s o l u t i o n to y i e l d formate ions. E i t h e r of these steps can be rate determining. I t would be expected that mass transf e r i s rate c o n t r o l l i n g under conditions of low a g i t a t i o n or where the rate of the second.step i s very f a s t , such as high temperature. A logarithmic p l o t of i n i t i a l rate against the inverse absolute temperature (Figure 12) shows t h i s to be indeed the case. Above about 90°C the r a t e -c o n t r o l l i n g step i s d i s s o l u t i o n of CO into the l i q u i d phase, as evidenced by the lower slope i n the p l o t . The k i n e t i c s of the formate s a l t formation • were studied most extensively at 80°C and t h i s i s w e l l within the reaction co n t r o l l e d region, where the slope of the p l o t i n Figure 12 i s much higher. A. Reaction C o n t r o l l e d Region  I d e n t i f i c a t i o n of Products I d e n t i f i c a t i o n of the reaction product was made from the buf f e r point of the f i n a l s o l u t i o n during t i t r a t i o n with standard HC1 which corresponded to the i o n i z a t i o n constant of formic a c i d (Figure 7). From the check of stoichiometry t h i s appears to be the only r e a c t i o n product that i s present i n appreciable concentration. Equivalents of HC1 (X 10 3) Figure 7. r. i ' . .iration Curve of F i n a l Formate Solution of Run Number 9 1 1 - 16 -• Stoichiometry of Reaction The measured stoichiometry was i n agreement with the reaction: CO + OH" — » - CH0 2" (4) The number of moles of OH" consumed i n the rea c t i o n was estimated from the analysis of the pressure-time curve i n Run number 9 and the r e s i d u a l OH concentration was obtained by di f f e r e n c e . The r e s i d u a l OH concentration was then experimentally determined' by t i t r a t i o n to pH 7 of the f i n a l s o l u t i o n from the autoclave (Figure 7). The formate concentration was determined by t i t r a t i n g f u r t h e r to a pH of 2.75 a f t e r the f i r s t end point was recorded. The r e s u l t s of stoichiometry checks in.Run number 9 a r e summarized i n Table I I . TABLE I I . V e r i f i c a t i o n of Stoichiometry of Run Number 9-Extent of Reaction x(moles CO) 0.0091 moles Residual OH" ([0H"] i-x) 0.0119 moles Analyzed [0H~] 0.0116 moles Deviation +0.000J moles Extent of Reaction (moles OH") analyzed 0.009^ moles Deviation -0.0003 moles Analyzed Formate O.OO9V3 moles Buffer Point (pH) 3.65 The end point of the formate t i t r a t i o n was cal c u l a t e d from the i o n i z a t i o n constant of formic a c i d and extent of reaction [H +][CH0P"] = 10~3.75 (5) [CH0 2H J or [ H + ] [ C H Q G " ] V l = l O " 5 ' 7 5 (6) extent of rea c t i o n - 17 -At the end point [H +] = [ C H 0 2 _ ] (7) Hence [H +] MlO'^'^ X extent of reaction (8) • v l This agrees w e l l with the end point estimated from the maximum slope of the t i t r a t i o n curve (Figure 7). K i n e t i c s The stoichiometry of the reaction indicates a bimolecular second-order reaction. The rate expression f o r t h i s i s written: . d [ C O t o t ] = . d[OH~ ] = d [ C H 0 2 - ] = k • [ C O d i s s ] [ O H _ ] ( 9 ) dt dt dt r where [CO^Q^] i s the t o t a l moles of CO divided by the l i q u i d volume at room temperature (83.4 cc) and [ C 0 ^ i s s ] i s the concentration of dissolved CO. Assuming Henry's Law to hold i n the region of pressures used (up to 35 atmospheres) [ C O d i s s ] = C X C 0 P C 0 ( 1 0 ) with()£,Q:== Henry's Law C o e f f i c i e n t At constant temperature and volume themples of CO i s proportional to pressure Nco = c p co ^ - 18 The number of moles of CO as a function of [OH -] i s given by ( W c o ) l - ([OH"]i -[OH"]) = [ C 0 t o f c ] ± - [0E-)± + [OH"] 2 V l where V]_ i s the l i q u i d volume i n the autoclave, equal to Qj>-.k cc at room temperature f o r a l l experiments. S u b s t i t u t i o n of equations(10),.(II) and(12) i n t o (9) gives: - d [ 2 0 = kO< V-L [CO,.]. - [OH -]. + [OH -] ) [0H _] dt : v "COT/ 1 1 C (13) Integration f o r non-stoichiometric (•[C0 t o t] 1 ^ [0H~] i) conditions gives: m. [cotot], - [ O H - ] . + [ Q H ~ ] _ l n [cotot l , = k O c ^ . - _ [ 0 H ] _ [OH -] [0H-] ± c 1 ( i M For the stoichiometric case ([CO^Q-t.^ = [OH"*"^) ToiPT T o T T . " k C * c o x t ( 1 5 ) 1 c Both C and <X are functions of temperature and C i s a function of l i q u i d 3 . volume... Values ofCX^Q were taken from the data of McAndrew . C was the slope of the experimentally determined c a l i b r a t i o n l i n e s described i n an e a r l i e r s e c t i o n . The value of C as a function of temperature, with a constant l i q u i d volume of 83 A cc at room temperature, i s given i n Figure 8. The use of the equation of state of CO at high temperature and pressure was not necessary with t h i s c a l i b r a t i o n . The experimental values of C and 0( are summarized i n Table I I I . Temperature (°C) Figure 8. V a r i a t i o n of Contained Gas Content C with Temperature - 20 -TABLE I I I . V a r i a t i o n of S o l u b i l i t y of Carbon Monoxide and Conversion Constants with Temperature Temperature °C. °<CQ ( a f t e r Mc Andrew) C (experimental) • X 10 4 moles atm'l'-70 6.7 9-65 80 6.7 9.23 90 6.8 8.82 100 7-1 8.40 The integrated rate law (14) was confirmed by a s e r i e s - o f experiments with various stoichiometric r e l a t i o n s between CO and. OH" as summarized, with the r e s u l t i n g rate constants, in.Table IV. TABLE IV. Experimental Rate Constants at 80°C. Experiment [K0H] i [ C 0 t o t J . Stoichiometry k X 10 2 Number M M 1 Relation M - 1 s e c _ 1 1 0.290 0.290 stoichiometric 2.62 £ , 2 0.290 0.144 C0-< 2.59 A 3 0.290 0.438 CO > 2.88 & 4 0.202 0.290 0H"'<" 2.48 ft 5 0.420 0.293 0H"> 2.59 a 6 0.285 O.7.29 CO -> 2.47 A Average - low group 2.61 * 10$ 7 0.168 0.221 0H"< 3.24 8 O.336 0.242 0H"> 3.28 9 0.252 0.114 CO < 3-77 . 10 0.252 0.342 CO -> 3.42 11 0.252 0.125 CO < 3-84 Average 3.51 ± 10$ The runs were performed with a f a u l t y technique which i s believed,to have resu l t e d i n c o n s i s t e n t l y low l i q u i d volume but these are consistent within.the group. - .21 -The rate constants were determined from the slopes of pl o t s such as Figure 9- The average rate constant f o r f i v e experiments i s 3.51 X 10~2 * 10fo M"1 s e c " 1 . .The two high values (Experiment number 9 and 11) correspond to experiments with the CO pressure l e s s than s t o i c h i o -metric. Mass tr a n s f e r c o n t r o l of the reaction i s more l i k e l y under low pressures and therefore the high values f o r these experiments p o s s i b l y r e f l e c t a mixed mechanism. I f these high rate constants are neglected the average rate constant i s 3-31 X 10" M - 1sec" . The deviation from l i n e a r i t y f o r the f i r s t f i v e minutes r e f l e c t s an i n i t i a l mass t r a n s f e r c o n t r o l of reaction rate, before the solutions were saturated with CO. Deviation from l i n e a r i t y at the higher times i s from s l i g h t l y i n c o r r e c t values chosen f o r equilibrium pressures. This does not appreciably e f f e c t the slope of the p l o t s during the ea r l y stages where they are l i n e a r . A pressure drop curve, c a l c u l a t e d from the average rate constant, i s shown i n Figure 6. A summary of a l l experimental data can be found i n Appendix B. .Effec t of Formate The e f f e c t of added potassium formate (0.213'M) w&s investigated i n Experiment number 12. No e f f e c t was noted (K^ = J>.2h X 10"2 M" 1sec" 1), thus the p o s s i b i l i t y of a pre-equilibrium i n the reaction mechanism was eliminated. S i r o t k i n 2 ^ found that CH02Na to a mole r a t i o of 1:1 with NaOH increased the rate of reaction. I t could be that his conditions are such that the mass tr a n s f e r of CO was the rate determining step. Formate ion could increase the reaction rate i n t h i s region, i f i t reduces the surface - 22 -Time (min) Figure 9. T y p i c a l Second Order Rate Plots - 23 -tension or the v i s c o s i t y of the solu t i o n . This p o s s i b i l i t y was not checked. E f f e c t of Cation NaOH and LiOH were substituted f o r KOH i n the reaction s o l u t i o n . The r e s u l t s are summarized i n Table V. TABLE V. . E f f e c t of Change of A l k a l i Metal Cation Experiment Number • Base Concentration M k r M" X 10 2 " 1 s e c _ 1 7 KOH 0.168 3. .2k 13 NaOH 0,200 5. • 25 Ik LiOH 0.230 3. .08 I f the lower rate i n LiOH i s a. r e a l e f f e c t i t r e f l e c t s the higher degree of asso c i a t i o n of l i t h i u m compounds, due to the much smaller i o n i c + radius of L i . To ca l c u l a t e the rate constant, i t was assumed the free hydroxyl ion concentration was the t o t a l LiOH concentration. A lower rate constant would be expected i f some as s o c i a t i o n existed, p o s s i b l y as ion p a i r s . However the above r e s u l t s are within the l i m i t s of precision expected f o r r e p r o d u c i b i l i t y and therefore do not require such an explanation. E f f e c t of Ionic Strength The i o n i c strength e f f e c t was investigated using 0.5 M, 1 M, 2 M NaC10 4 A. A NaC104 was chosen i n preference to KC104. due to i t s greater s o l u b i l i t y . I t was shown i n previous section,Na + i n place of K has no e f f e c t . -24 -The e f f e c t of increasing i o n i c strength was to decrease the rate constant. 1 M Na 2S0 4 also decreased the rate constant. C o r r e c t i o n was made i n the sulphate system f o r the decrease i n CO s o l u b i l i t y as found by Forward . The r e s u l t s are summarized i n Table VI. TABLE VI. E f f e c t of Ionic Strength Experiment NaOH NaC10 4 O<C0 A _n k r X 102 Number M M M ; • tatm x M " i s e c - 1 U 0.200 6.7 X 10" 4 3.25 15 0.200 0.50 6.7 3.02 16 0.200 1.00 6.7 2.85 17 0.200 2.02 6.7 2.48 18 Na 2S0 4 - 4 0.200 1.00 4.5 X 10 2.61 E f f e c t of Catalysts The e f f e c t s of several ions that might catalyze t h i s reaction were investij gated. The r e s u l t s are summarized i n Table VII. TABLE VII. .Effects of Ions as Catalysts on Rate of Formation of Formate S a l t s Experiment K0H Added Salt Concentration k X 10 2 Number M M - 1 s e c _ 1 7 0.168 3.24 19 0.252 KMn04 0.0095 3.24 20 0.252 KNO3 0.011 3.26 21 0.252 T l N0 3 O.OO87 3.21 22 0.252 Ag(EDA) 2 CIO4 0.013 3.24 • 23 0.252 Cu(EDA) 2(C10 4) 2 0.0095 2.66-4.07 - 25 -No e f f e c t from these ions was noted. Cu(C10 4) 2 and AgC10 4 were added as ethylene diamine complexes to prevent hydrolysis i n the strongly a l k a l i n e solutions. Interpretation of the Cu(II) experiment was made d i f f i c u l t , by the CO-reduction of Cu to metal which was competing with the OH - ions f o r CO. I f t h i s experiment i s analyzed on the assumption that no Cu i s reduced a value of 2.66 X 10~ 2 M - 1 s e c - 1 i s obtained f o r the rate constant. When the Cu reduction to metal, according to the. stoichiometry Cu(II) + CO + 20H" — > - Cu° + H 2C0 3 '.(16) was assumed complete before the formate reaction began, the rate constant calculated from the pressure drop curve was 4.07 X 10- 2M _ 1sec - 1. Both of these values are not f a r from the average value f o r the formate s a l t formation and any c a t a l y t i c e f f e c t i s small. Halpern detected the presence of formate i n the f i n a l solutions • 27 of the s i l v e r reduction by CO . The a n a l y t i c a l technique was not described, but an estimate of the amount' of formate produced, under Nakamura-'s 2 experimental conditions was made on the basis of the present reaction k i n e t i c s and indicated about 1 0 - 7 M formate would be formed i n a t y p i c a l experiment. Detection of t h i s concentration i s doubtful and there seems to be some evidence f o r a c a t a l y t i c e f f e c t from about equimolar concentrations of A g + ions, although no e f f e c t was found from the concentration used i n t h i s work. - 26 -Rate i n Buffer Solution The reaction was invest i g a t e d - i n a KHPO4/K3PO4 b u f f e r system. The rate law i n a buffer system becomes: - d [ C ° t o t ] B k r [ 0 H - ] [ C 0 d i s s ] ' ( 1 7 ) dt w i t h ' [ O H ] almost constant. The dissolved CO i s rela t e d to the t o t a l CO by: [ c o d i s s ] = ° < c o p co ' ( l 8 a ) -• = ^ c o N c o (18b) c = O < C 0 V l _ . [ C 0 t 0 t ] ( 1 8 c ) C Equation (12) becomes -^[cotot] = ^ V j . [OH - ] [ c o t o t ] ( i 9 ) dt 7j The integrated form with [ O H ] considered constant i s - In [CO] + In [CO]^ = - k^Ccfe [0H~] t • . (20) C The slope (k). of a p l o t of - ln[C0] against time i s : ^ = -k rO< C 0 Vl :[0H"] (21) C ~ The hydroxyl ion concentration was calculated from the slope, assuming that the rate constant was the same i n t h i s buffered system. Ionic strength e f f e c t s on the rate constant have already been.shown.to be small,, as were the c a t a l y t i c e f f e c t s of various ions. - 27 -.The t h i r d i o n i z a t i o n constant was calculated from the hydroxyl ion concentrations. - For.the reaction HP0 4 - + 0H~ P 0 4 ^ + H £0 (22) the equilibrium constant i s K d = [F0 4= ] (23) L H i P 4 = J l O H _ J . The t h i r d i o n i z a t i o n constant f o r phosphoric a c i d i s ( K i L = [ P 0 4 s ] [ H + ] ,,,-.(24) . [ H P O 4-J From equation (rf)and (18) ( K i ) 3 = K a K y (2 5) where K i s the i o n i z a t i o n constant of pure water. Thus w ( K i ) 3 = [P0 4^] K w (26) ,[HP04"][OH"] The experiment (number 24) was done i n a-1.74/1 K3P04/K2HP04. buffer s o l u t i o n of t o t a l phosphate concentration equal to 2.68 M at 80°C. .The r e s u l t s are shown i n Figure 10, the slope corresponds to a hydroxyl ft ion concentration of.3.59 'X 10"3.M ft The r a t i o of Henry's-Law, C o e f f i c i e n t s f o r CO i n water and phosphate buff e r s o l u t i o n was determined at room temperature to be 3-2..It -was assumed that t h i s r a t i o was the same at 80°C. hence ®{QQ i n phosphate =2.11 10~4 M atmos" 1. - 28 -•1550 h .1500L .1^ 50 I -p o -p o o 3 .1400 .1350 .1320 h 0 T o 60" Time (min) 80 100 Figure 10. F i r s t Order Rate Plot i n Phosphate Buffer - 29 The amount of phosphate di s s o c i a t e d to give the hydroxyl ions was within the range of e r r o r of the c a l c u l a t i o n and neglected. 28 The i o n i z a t i o n constant of water i s given by Harned and Owen as l o g K w = -M7O.90 + 6.O875 - .O.706O T ....,(27) At 80°C K^ . i s 2.60-X 10" 1 3. The t h i r d i o n i z a t i o n constant of phosphoric a c i d at 80°C was calculated to be (K ) = 1.26 X 10" 1 0 from equation (26). i 3 v 28 This i s about a 250-fold increase from the 25°C value of 4.8 X 10" 1 3 The low temperature value i s f o r a zero i o n i c strength solution.whereas the i o n i c strength of the-buffer s o l u t i o n used i n the run was about ten. This i s much, higher than the-Debye-Huckel range and no acceptable theory of i o n i c strength e f f e c t s at t h i s high concentration e x i s t s . Although the i o n i c strength e f f e c t on the rate constant of the formate reaction was small there i s c e r t a i n to be an e f f e c t on the i o n i z a t i o n constants of water and phosphoric a c i d . A An estimate of the i o n i c strength e f f e c t gave an i o n i z a t i o n constant of 1.26 X 10~9 at 80°C. .4 The pH of the buffer s o l u t i o n at room temperature was measured with a glass electrode standardized with a.3.55X 10"2 M .NaOH so l u t i o n . From the l i t e r a t u r e value of the i o n i z a t i o n constant the pH of each of these solutions should be the same neglecting i o n i c strength. A pH of one u n i t higher was found f o r the buffer s o l u t i o n and was a t t r i b u t e d to the i o n i c strength of the s o l u t i o n . - 30 The average heat of i o n i z a t i o n of mohohydrogen phosphate was calculated to be 25.5 ± 5 Kcal/mole from the average slope of a ;plot of log K3; against the inverse absolute temperature (Figure 11). I t i s u n l i k e l y that the heat of i o n i z a t i o n i s constant over the temperature range studied, the v a r i a t i o n of the heat of i o n i z a t i o n with temperature i s given by K i r c h o f f ' s Law, T 2 \ r e l a t e s t h i s change to changes i n s p e c i f i c heat A g . The s p e c i f i c heat of s a l t s i n solu t i o n .is d i f f i c u l t to obtain because of solvent e f f e c t s Two opposing e f f e c t s influence the d i s s o c i a t i o n of an a c i d with temperature and are reflected-, i n the heat of i o n i z a t i o n . -The d i s s o c i a t i o n of the a c i d required a heat of d i s s o c i a t i o n , and the solva t i o n of the r e s u l t i n g ions gives a heat of sol v a t i o n . C e r t a i n l y s o l v a t i o n of a sym-metric ion such as P 0 4 = w i l l be an exothermic process, r e l a t i v e to the HPO4 ion. D i s s o c i a t i o n w i l l generally be an endothermic::reaction and more important at higher.temperatures. • The high charge on the H P O 4 ion should r e s u l t i n a high heat of d i s s o c i a t i o n . The e f f e c t of i o n i c strength i s to make both processes more exothermic. The heat of i o n i z a t i o n i s the sum of these enthalpies and u n l i k e l y to be constant over the-whole temp-erature range. - The above c a l c u l a t i o n s i l l u s t r a t e the p o s s i b i l i t y of obtaining thermodynamic data from the k i n e t i c s of the formate s a l t formation reaction The i o n i z a t i o n constants and heat of i o n i z a t i o n of weak acids and bases, such as ammines, some inorganic bases,.and polybasic acids, can be obtained Figure 11. Heat of Ionization of Monohydrogen Phosphate -32 -at elevated temperatures - conditions f o r ' which data i s lacking. Several c r i t e r i a must be f u l f i l l e d . t o make t h i s technique u s e f u l , these beings a. the acid- or base must be i n e r t to reduction or reaction with carbon monoxide. b. the system must be buffered. However, consistent with t h i s the i o n i c strength should be kept as small as p o s s i b l e , to minimize complications. c. the rate of reaction must be slow enough to prevent mass tr a n s f e r c o n t r o l of the formate formation as t h i s i s • independent of pH. • The l a s t c r i t e r i o n i s not too c r i t i c a l as the* reaction may be slowed down by proper reagent c o n t r o l . Decreasing the concentrations decreases the rate. Using other methods•of analysis of the reaction rate, such as analysis of samples, or by continuous spectrophotometric observation of the co l o r change of a suitable i n d i c a t o r , very slow re a c t i o n rates may possibly be followed. . E f f e c t of Temperature The r e a c t i o n was studied.over the temperature range of 70°C to 200°C. -An Arrhenius p l o t of the i n i t i a l rate ( _ d?QQ ) against the - d T -inverse absolute temperature i s shown i n Figure 12. There i s a change of slope around 90°C .corresponding to the conditions where the reaction becomes d i f f u s i o n c o n t r o l l e d . This w i l l be discussed i n a l a t e r section. Figure 12. Arrhenius Plot of I n i t i a l Reaction Rate TABLE VIII. • Data f o r Arrhenius Plot Experiment Temperature k r X 102 , I n i t i a l Rate Number • • °C. • M - 1 s e c - 1 X' 10 2 atmcgec-i 26 70 •1.10 0.431 27 80 2.55 1.12 28 90 6.58 3.19 29 100 I6.5O 9.05 k i : X 10 3 a s e c - i 29a 100 1.43 4.07 30 110 1.93 5.44 31 120 2.59 7-38 32 150 4.72 14.2 33 200 16.80 •47.9 a). A c t i v a t i o n - Energy In the reaction controlled, (low. temperature) region the slope of the l i n e corresponds t o an a c t i v a t i o n energy of 23.2 ± 1.5 Kcal.mole - 1. 21 22 This i s i n good agreement with the r e s u l t s of Kodama ' . The dif f e r e n c e i n activation.energy f o r Na + and N H 4 + reported by Kodama i s p a r t i a l l y accounted f o r by the heat of i o n i z a t i o n of ammonium hydroxide. The heat of i o n i z a t i o n of N H 4 O H was calculated from the slope of log (K_. ) ^ 1 JN.n40ii against the inverse absolute temperature to be 2 * 0.2 kcal/mole. b) Entropy of A c t i v a t i o n The entropy of a c t i v a t i o n was calculated from the equation (29): _ E a A S * kr. = e K T "e R T e ~^ (29) h and. the experimental data at 80°C,to be A s ^ -2.2 ± 1 eu. The normal entropy change f o r a bimolecular a s s o c i a t i o n .reaction i n the gas phase i s about -10 eu. In the l i q u i d phase entropy changes i n the solvent a l s o occur. -Coordination of water molecules about a formate ion i s much more -35 -random than about r e l a t i v e l y small ions l i k e OH and.CO;. -This would account f o r some of the high entropy change, r e f l e c t i n g the decrease i n order i n the solvent. An increase i n As"^ can also r e f l e c t an increase i n rea c t i o n cross-section. 30 Proton t r a n s f e r in.water i s very f a s t ^ and therefore i t is' possible that an OH ion i n the v i c i n i t y of a CO molecule could cause i o n i z a t i o n of a water:molecule adjacent to the CO, as C ^ /H .H (30) 0 JS. 0 - H X H .0 _ H — - (T The CO molecule could also a i d by f o r c i n g the water molecule to io n i z e . I f this accounts f o r part of the small entropy decrease of t h i s reaction, i t i s evident only a-minority of water molecules are oriented with OH ions to p a r t i c i p a t e t h i s way. Otherwise dependence of the rea c t i o n rate on-OH" concentration would disappear, because the r e a c t i o n - i s e s s e n t i a l l y with water molecules rather than OH ions. - The Reaction Mechanism A mechanism consistent with the above discussed r e s u l t s can be described by the following equation: _ k r CO -+.0H — > . ' [CO - OH ] (31a) f a s t // ° [CO - OH"] — 3 - H - C (31b) 0._ - 36 -This r e a c t i o n scheme gives r i s e t o t h e rate expression d I C O t a t ] _ -d[OH~] _ d.[CO-OH"] = k r [CO ]. [0H~] (32) dt dt dt The absence of a c a t a l y t i c e f f e c t would indi c a t e a straight-forward mechanism.and the k i n e t i c s were i n agreement with t h i s rate law. The retardation of the reaction.with i n c r e a s i n g . i o n i c strength i s evidence f o r ions of opposite charge taking part. The magnitude of this, e f f e c t was smaller than would be expected from ions-of unit change. This indicates the carbon atom of the s l i g h t l y d i p o l a r (0.1 D^) carbon monoxide molecule being the reactive end since i t i s the p o s i t i v e end of the dipole. The nature of the complex.(CO - OH ] (equation 31) i s not defined above. Two possible configurations may be considered of these, the more reasonable, i n terms of bond breaking i s : ,o-\ OH A f a s t proton t r a n s f e r would follow, p o s s i b l y by way of solvent molecules. 3 - 0 H C C £ ^ 1 H-CT I ^ O H X 0 -0 X H .(33) / 0. 0 -H 0 H <H • XH / X H The a l t e r n a t i v e i s a complex much l i k e the- formate ion - 37 -•A f a s t rearrangement of the bonds then gives the formate ion ! ~-C = 0 .0'-f a s t , 0 +• H - C (34) 0 ^ . This i s an i n s e r t i o n complex s i m i l a r to those postulated f o r other CO reactions. The d i s t i n c t i o n between the above a l t e r n a t i v e s was interpreted from an experiment i n D 20 at 80°C. A s o l u t i o n of KOH (0.291M) was made up i n 99.5$ D 20. The rate constant found was 4.28 X 10~2 m o l e - 1 s e c - : L , assuming the s o l u b i l i t y of CO was the same i n D20 as H20. This represents an increase i n rate and i s equivalent to a separation.factor ^ - f e 0 / 7 7 <»> . The increased rate i n D 20 rules out the p o s s i b i l i t y of proton t r a n s f e r i n the rate determining! step; deutron transfer.- rates are always slower than rates f o r protons. 32 According to Eyring and Cagle , a.value of l e s s than one i s a very u n l i k e l y occurrence unless two "conditions are f u l f i l l e d : i . the temperature of reaction must be i n the r e l a t i v e l y low region of quantum e f f e c t s , i . e . kT <iC h Q . i i . the sum of the h a l f quanta-of the v i b r a t i o n a l frequencies of the a c t i v a t e d complex, i n the d i r e c t i o n of the reaction ± S o l u b i l i t i e s of s a l t s generally decrease i n D 20. I f t h i s i s true a l s o f o r CO the e f f e c t would r e s u l t i n a further increase i n the rate constant. - 3 8 -coordinate,.must be greater.than the sum of the equivalent frequencies i n the reactants. The ground.state v i b r a t i o n a l s t retching frequency f o r the hydroxyl ion bond i s equivalent" to 2.3ytjC^ • The thermal energy at 80°C i s about one twentieth of the quantum energy hence the f i r s t of these conditions f o r a f r a c t i o n a l cX value is. s a t i s f i e d . The second.condition i s the basis of the zero-point energy 3 6 c a l c u l a t i o n as discussed by Melander . .This c a l c u l a t i o n gives strong evidence f o r the existence of the i n s e r t i o n complex in. the reaction mechanism. Table IX. l i s t s the wave numbers of the ground, state stretching modes of the indicated bond. TABLE.IX. Vibr a t i o n a l . Stretching Frequencies .Bond • Wave Number • (£j) cm -i Reference OH i n OH" 4 3 5 0 33 OH i n CH02H . 3 5 7 0 34 CH i n CH02H 2985 3 4 .•OD i n OD" 3 2 2 5 33 O^H i n CD02D 2 6 8 0 3 5 .CD i n CD02D 2265 35 From.these values the difference i n zero point energy f o r the hydroxyl ion i s A(ZFE) 0 H - . Q D = V 2 h c ^ H - - l / . h c ^ - ( 5 6 ) 1 . 6 1 X 1 0 3 c a l mole" 1 Formic a c i d was used as an approximation f o r the formate ion i n these c a l c u l a t i o n s . For the CH bond of formic a c i d : A ( Z P E ) C H ^ D = 2.65 X 103 cal/mole (37) I t was assumed:the change i n zero point energy of the activa t e d complex was the numerical average of the above values, c a l c u l a t e d f o r the reactant and product. The separation f a c t o r (o( ) was calculated from: CX = ^ = exp - (%20 - E D p Q ) (38) D20 RT The change i n a c t i v a t i o n energy (Eg^Q — E ^ Q ) i s given by: ( EH 20 - ED20> = - 1 / 2 ( A ( Z f £ ) 0 H - . 0 D - - ^ ( Z P E ) ^ ^ ) (39) from which .k. 0( = j M = M (4o) KD 20 An analogous c a l c u l a t i o n f o r the complex configuration r e q u i r i n g subsequent proton t r a n s f e r gave 0( =1.26 corresponding to a decreased rate i n D 20. Experimentally, as noted on page 37, 0{ i s 0.77- This i s i n reasonable agreement with the expected value i f the k i n e t i c e f f e c t of the increased hydrogen mass in.the solvent i s considered i . e . 0.77 1/ ^ D 2 ° = 0.77X0.71= O.55 (4l) f/^O w h e r e i s the reduced mass. 4o Thus the heavy water run.supports a mechanism.of formate ion formation described by the equations k r 3 « — 0' •CO^  . + OH d i s s i PC = 0 H--C = 0 f a s t H- C 0 \ « -(42) B. • D i f f u s i o n Controlled Region  K i n e t i c s For d i f f u s i o n processes the rate i s given by Fields Law dc dt K & d ? .(kk) In the d i f f u s i o n controlled.region of the reaction the r e s i d u a l CO concen-t r a t i o n i n the l i q u i d phase approaches zero. The concentration gradient across the g a s - l i q u i d i n t e r f a c e i s therefore proportional to the pressure and the rate equation becomes: -dP CO dt The. integrated form becomes CO In CO ( P c o ) i k d t .(46) where k d i s proportional t o the mass tr a n s f e r c o e f f i c i e n t . A p l o t of In PC0/( -^oo) • a S a i n s t t gave a.straight l i n e with slope equal • k^. T y p i c a l p l o t s are shown i n Figure 13. -As the reaction proceeds the rate slows down and becomes reaction c o n t r o l l e d , where i t changes from f i r s t order.kinetics and.deviates from a s t r a i g h t l i n e i n the f i r s t order p l o t . - hi -Time (min) Figure 1$. F i r s t Order Rate Plots i n Mass Transfer Controlled Region - 42 Second order plo t s were made (Figure 14) and rate constants calculated. .This approach was used to determine the rate constant (k ) at 100°C. The E f f e c t of Temperature The v a r i a t i o n of d i f f u s i o n c o e f f i c i e n t s with temperature i s shown i n Table X , as w e l l as the i n i t i a l rates used f o r the Arrhenius p l o t (Figure 12). - TABLE X . V a r i a t i o n of D i f f u s i o n C o e f f i c i e n t s and I n i t i a l D i f f u s i o n Rates with Temperature Experiment Temperature k^ I n i t i a l Rate Number ° C . . X -J_Q3 atmos s e c - 1 X 1 0 2 29a 1 0 0 1 . 4 3 4 . 0 7 3 0 110 1 . 9 3 5 . 4 4 3 1 1 2 0 2 . 5 9 7 . 3 8 32 150 4 . 7 2 1 4 . 2 33 2 0 0 1 6 . 8 4 7 . 9 The slope of the Arrhenius p l o t i n the high temperature region corresponds to an a c t i v a t i o n energy f o r mass t r a n s f e r of about 9 - 0 Kcal. This i s higher than normal experience f o r d i f f u s i o n i n the aqueous phase which i s u s u a l l y l e s s than 5 Kcal. The heat of s o l u t i o n f o r CO was not considered i n the above c a l -c u l a t i o n . -From the slope of a p l o t of log &\QQ against temperature the heat of s o l u t i o n was estimated to be 3 - 4 k c a l mole" 1. . Thisi.heat of s o l u t i o n accounts f o r most of the excess a c t i v a t i o n energy i n the d i f f u s i o n c o n t r o l l e d region. The area of the in t e r f a c e a l s o e f f e c t s the rate of d i f f u s i o n and consequently the apparent a c t i v a t i o n energy. The effectiveness of the a g i t a t i o n was decreased at higher reaction temperatures because the l i q u i d - kk -phase expanded, r a p i d l y decreasing the r e l a t i v e l y small gas'volume. This would decrease the surface area and r e s u l t i n a smaller apparent a c t i v a t i o n energy. This e f f e c t must be small, since factors, r a i s i n g the a c t i v a t i o n energy above that of simple d i f f u s i o n seem to dominate. The true mechanism of a homogeneous reaction i n s o l u t i o n i s obscured, i f mass t r a n s f e r i s the r a t e - c o n t r o l l i n g step. The l i m i t i n g rate that can be studied i n the shaking autoclave apparatus, from the mechanistic approach, was estimated from the mass tr a n s f e r k i n e t i c s . As a function of temperature t h i s maximum rate of pressure drop, f o r carbon monoxide, i s _ 9000 RT - a p c o = 2.51 x 10 2 x P c o e (47) dt C. .Applications of the Formate Reaction It has been shown that basic solutions are n e u t r a l i z e d by the absorption of carbon monoxide. • I f the formate ions produced are oxidized, mixtures of carbonate and bicarbonate form and the pH of solutions containing these ions can be r a i s e d t o about 12 by b o i l i n g and s t r i p p i n g out the carbon dioxide. This cycle could be applied to an i n d u s t r i a l process r e q u i r i n g successive lowering and r a i s i n g of pH l e v e l s , provided that the presence of formate or carbonate s a l t s ; i s not detrimental. The leaching of s i l i c a or alumina-ores i s a process that could be subjected to t h i s type of procedure. These minerals may be leached by basic s o l u t i o n and then p r e c i p i t a t e d by a c i d i f i c a t i o n . The a c i d i f i c a t i o n may be e f f e c t e d by passing e l e c t r i c furnace gas through the solution.- S i l i c a or alumina w i l l p r e c i p i t a t e a f t e r which a i r or oxygen, blown through the acid s o l u t i o n , oxidizes the formate to carbonate and more mineral may be dissolved and the cycle repeated. • Scheelite ores may be s i m i l a r l y treated, with the production of synthetic s c h e e l i t e (CaW04) by liming the formate solutions to remove tungsten before r e o x i d i z i n g to regenerate the sodium carbonate leach s o l u t i o n needed to dissolve s c h e e l i t e . he -CONCLUSION The absorption of carbon monoxide by basic s o l u t i o n i s described by the reaction CO + OH" =~ CH0 2 _ Added reagents appear to have no e f f e c t on the k i n e t i c s of the reaction, although an increased i o n i c strength exhibits a s l i g h t retardation e f f e c t . A pronounced increase i n reaction rate occurs i n deuterium oxide. A mechanism consistent with the observed k i n e t i c s i s described by CO + OH ;c = o sc = o f a s t i / • 0" H - C \ , 0 The absence of an e f f e c t from added reagents allows the i o n i z a t i o n constants of weak acids and bases to be determined from the k i n e t i c s of the formate s a l t formation. Mass tr a n s f e r of carbon monoxide from the gas phase to the l i q u i d phase becomes rate c o n t r o l l i n g f o r the conditions used i n t h i s study, when the rate of the homogeneous reaction i s greater than - d PC0 = 2.51 X 10 2 PC0 6 dt 9000 RT I n i t i a l rates exceeded t h i s i f the reaction temperature was greater than 90°C . APPENDIX A. Observations Concerning Reduction of Cobalt(II) by Carbon Monoxide P r i o r to the study of the formate s a l t formation an attempt to study the k i n e t i c s of a possible reduction of cobalt with carbon monoxide was undertaken. Two experimental techniques were t r i e d . The f i r s t of these was using the constant volume shaking autoclave already described, while the second was with a high pressure spectrophotometric c e l l as described 37 by Byerley and Peters •Several runs were done i n the shaking autoclave i n an attempt to f i n d conditions under which the reaction rate was s i g n i f i c a n t . C0SO4 concentration of about 0.01 M i n (NH 4) 2S0 4/NH 40H buffer s o l u t i o n and carbon monoxide p a r t i a l pressures of 700 p s i g were used. Temperature was varied from 100 to 220°C. • A s l i g h t pressure drop, was observed between CO and 0 2 i n pure water, the reaction with Co(II) was studied with a gas mixture of CO and 0 2 (4/l C0/0 2) at 200°C. Again a.very slow pressure drop was noted. • A slow pressure drop. .(indicating a.reaction) was also observed i n a (NH 4) 2S0 4/NH 40H buf f e r s o l u t i o n . A run i n unbuffered s o l u t i o n indicated no reaction. It can now be understood that the slow pressure drop was due to the carbon monoxide reacting with the buffer to give formate ions. The pressure drop technique was abandoned a f t e r these runs. I t was attempted to follow the reduction of cobalt s p e c t r o s c o p i c a l l y . Ethylene diamine complexes of cobalt were used i n the c e l l . This prevented hydrolysis ++ of Co(II) i n basic s o l u t i o n . 0.01 M Co(EDA)3 solutions were put i n the +++ spectrophotometric c e l l and heated to 120°C. The reduction of Co(EDA) 3 was checked i n both N 2 and CO. No decrease of absorbance resulted. - 48 -A slow decrease i n absorbance occurred i n (NH 4) 2S0 4/NH 4OH ++ buffered s o l u t i o n of Co(NH 3) 6 with CO at 120°C. No reaction was detected below t h i s temperature. -On heating the ammonia buffered so l u t i o n to 150°C under N 2 the windows become coated with a p r e c i p i t a t e . I f CO i s introduced the p r e c i p i t a t e redissolves and there i s a rapid decrease i n absorbance. No reduction i s detected i n the acetate buffered system (pHss5). No s a t i s f a c t o r y explanation of these e f f e c t s i s given. - I t could be that there i s a thermal decomposition process occuring. I t does appear that carbon monoxide or formate ions are s o l u b i l i z i n g the p r e c i p i t a t e i n the high temperature runs. A d d i t i o n a l Spectrophotometric Runs 2 + The reduction of the U0 2 ion by CO was investigated. No d i r e c t reduction was detected at 120°C. No c a t a l y t i c e f f e c t of Cu(II) or Cu(CO) + was detected,upvto a pH 5 with equi^molar U 0 2 2 + and Cu +(CO). 2 + The reduction of Ni with CO was detected spectrophotometrically at temperatures up to l60°C i n buffered solutions. This rate was however slow and no evidence of a Cu(II)-Cu(I) c a t a l y t i c e f f e c t was detected. No c a t a l y t i c e f f e c t of N i 2 + toward U 0 2 2 + was evidenced. 38 Bawn and-White studied the reduction of Co(III) to Co(II) with formic a c i d and concluded that there was a two path mechanism: Co(III) + HC02H — » - Co.(II). + HC0 2 _ + H + (48a) Co(III). + HC0 2" —*- Co(II). + HC0 2 _ .(48b) followed by f a s t reactions of the formate r a d i c a l s formed. - K9 -Perhaps under more d r a s t i c conditions of temperature Co(II) could be reduced by formic a c i d . I t i s expected the reduction of Co(II) i s i n v e r s e l y dependent on the hydrogen ion concentration. By s t a r t i n g with formic acid, instead of CO, higher pH solutions could be used without complications from CO reaction to give formates. APPENDIX B. Summary of Experimental Data - 50 -Experiment I n i t i a l .CO I n i t i a l Temp. k r (X 102) Number Pressure KOH Con-: °C. Remarks X 10 4 (atm) centration M atm - 1 M~ 1sec- 1 M 1 2 6 . 2 2 0 . 2 9 0 8 0 2 1 2 . 9 6 0 . 2 9 0 8 0 3 3 9 - 5 5 0 . 2 9 0 8 0 4 2 6 . 1 9 0 . 2 0 2 8 0 5 2 6 A 8 0 . 4 2 0 8 0 6 6 5 . 6 0 O .285 8 0 7 1 9 . 9 ^ 0 . 1 6 8 8 0 8 2 1 . 9 0 0 . 3 3 6 8 0 9 1 0 . 2 7 0 . 2 5 2 8 0 10 3 0 . 7 4 0 . 2 5 2 8 0 11 1 1 . 0 6 0 . 2 5 2 8 0 12 3 0 . 9 0 0 . 2 5 2 8 0 13 3 0 . 9 2 8 0 14 3 0 . 7 0 8 0 15 3 1 . 0 9 0 . 2 0 0 (NaOH) 8 0 16 3 0 . 7 6 0 . , 2 0 0 (NaOH) 8 0 17 3 1 . 2 8 0 , , 2 0 0 (NaOH) 8 0 18 3 1 . 1 3 0 . , 2 0 0 (NaOH) 8 0 19 3 0 . 8 9 0 . 2 5 2 8 0 2 0 3 1 . 5 2 0 . 2 5 2 8 0 2 1 3 0 . 9 7 0 . 2 5 2 8 0 22 3 0 . 8 9 0 . 2 5 2 8 0 23 3 0 . 7 3 0 . 2 5 2 8 0 2 3 a 2 9 . 3 6 0 . 1 9 2 8 0 2k 2 4 . 1 3 8 0 25 3 0 . 7 2 0 . 2 9 1 8 0 26 2 4 . 7 6 0 . 2 7 3 7 0 27 2 6 . 0 0 0 . 2 7 3 8 0 2 8 2 6 . 0 6 0 . 2 7 3 9 0 29 2 8 . 5 1 0 . 2 7 3 100 29a 2 8 . 5 1 0 . 2 7 3 1 0 0 3 0 3 1 . 1 5 0 . 2 7 3 110 3 1 3 1 . 5 5 0 . 2 7 3 1 2 0 32 3 5 - 6 6 0 . 2 7 3 150 33 4 1 . 2 7 0 . 2 7 3 2 0 0 6 . 7 2 . 6 2 6 . 7 2 • 59 6 . 7 2 . 8 8 6 . 7 2 .48 6 . 7 2 • 59 6 . 7 2 • 47 6 . 7 3 .24 6 . 7 3 . 2 8 6 . 7 3 .84 6 . 7 5-.42 6 . 7 3 • 96 0 . 2 1 3 M-KCHO2 6 . 7 3 .24 0 . 2 0 0 M-NaOH 6 . 7 3 • 25 0 . 2 3 0 M-LiOH 6 . 7 3 . 0 8 O.5OO M-NaC104 6 . 7 3 . 0 2 1 . 0 M-NaC104 6 . 7 2 . 8 5 2 . 0 2 M-NaC104 6 . 7 2 .48 1 . 0 M-Na 2S0 4 ^ • 5 2 . 6 1 0 . 0 0 9 4 6 M-KMn04 6 . 7 3 .24 0 . 0 1 1 1 M-KNO3 6 . 7 3 . 3 6 0 . 0 0 8 6 8 M-TINO3 6 . 7 3 . 2 1 O.OO999 M 6 . 7 3 .24 Ag(EDA) 2C10 4 0 . 0 1 2 4 5 M 6 . 7 2, .66 Cu(EDA)(C10 4) 2 0 . 0 1 2 4 5 M 6 . 7 4, , 0 9 Cu(EDA)(C10 4) 2 1 . 2 0 M-K 3P0 4 2 . 1 O .98 M-K2HP04 9 9 - 5 $ D 2 0 6 . 7 4 . 2 8 6 . 7 l . 1 0 6 . 7 2 .55 6 . 8 6, . 5 8 7 - 1 16, .5 k d ( x 1 0 f ) s e c " 1 1 4 3 1 . 9 3 2 . 5 9 4 . 7 2 1 6 . 8 A Copper reduction assumed complete before formate reaction began. 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Melander, Isotope E f f e c t s on Reaction Mechanisms, Ronald Press Company, New York ( i 9 6 0 ) . 3 7 - E. Peters and J.. J . Byerley, Rev. of S c i . Inst. 3jt> 819 ( I 9 6 3 ) . 3 8 . C.E.H. Bawn and A. G.. White, J.. Chem. Soc. 3 3 9 , ( 1 9 5 1 ) . 

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