This dissertation has been microfilmed exactly as received Mic 61-929

MERRYMAN, Earl Lewis. THE ISOTOPIC EXCHANGE REACTION BETWEEN Mn AND MnO” . 4

The Ohio State University, Ph.D, 1960 Chemistry, physical

University Microfilms, Inc., Ann Arbor, Michigan THE ISOTOPIC EXCHANGE REACTION

BETTAIEEN Mn** AND ItaO ^

DISSERTATION

Presented in P&rtial Fulfillment of the Requirements for the Degree Doctor of Philosophy In the Graduate School of The Ohio S tate U niversity

By

Earl Lewis Ferryman, B.Sc*

The Ohio State University

I960

Approved by

Department oy Chenletry 1C mnriEDGiBiT

The author wlshea to e:qpr«as his approoiation to

Profoaaor Alfred B. Garrett for hie superrieion and enocur- agement during the oouree of this research* and for his sincere interest in mj eelfare both as an undergraduate and graduate student at Ohio State University. I also wish to thank the Ohio

State University Cheidstry Depsurtnent for the Assistant ships granted me during the 1 9 5 6* 7 "^ aeademlo years.

The author also gratefully acknowledges the Fellowships granted me by the American Cyansuald Company during the 1959*60 academic year and by the National Science Foundation during the

Summer Q u a rte r of I960*

i i TABI£ OP CONTEHTS PAOE

INTRODUCTION ...... 1

Àpplloationa of Radloaotirlty in Chomiatry 1

The Problem and Its H latory ...... 1

The Problem Reeulting from Early Work 5

Statement of the Problem ...... 7

Stability of Varioue Oxidation States of Mengeneee •••..• 9

MançaneeeCl) ...... 9

Wanganeae(ll) * ...... 10

Manga ne ee (III) ...... 10

Manganese(IV) ...... »...... 11

Manganese(V)« (Vl) and (VIl) ...... 12

EXPERIICNTAL HiOCEDURES ...... 15

Preparation and Purification of Reagents ...... 15

Potassium Permanganate Solution ...... 15

Manganous Sulfate Solution 15

Procedure Used for Recovery of Manganous Ions ...... 17

Procedure for Determining Permanganate Using Iodide Ion . 18

Preparation of Radioactive Manganese, Ifn^^ ...... 18

Iron Target Bombarded with Deuterons, (12 Mb t.) ...... 18

Experimental Prooedure for Studying the Isotopic Exchange Reaction Between Mn*^ and iVsO]^ ...... 23

Experimental Prooedure to Determine the Time of Formation of Manganese Dioxide for Studying the K in e tic s of th e Mi -MnO^ S y s te m 24

Experimental Prooedure to Determine whether Manganous Ions Form Complexes ...... 25

i l l T1B1£ OF OONTEBTS (oontijaued) PiOE

PrapATfttlon of Anion Exehango Column# ...... 26

THEORY ...... 28

EXFERIlffiHTAL DATA AND RESULTS ON RECOVERY OF lUNGANESE...... 31

Molarity of Solutions Used in the Experiments ...... 31

R eooeery D ata on Perm anganate S o lu tio n s ...... 31

Conclusions from the Recovery Data in Tables 1 ,2 and 3* 36

Recovery Data from Solutions and Mn*^*-MaOT Solutions ...... 38

DATA AND RESULTS FROM THE ISOTOPIC EXCHANOE EXPERIMENTS 46

Discussion of the Results Obtained from the Mn'^^-MnO]^ Isotopic Exchange Reaction ...... 50

DETERMINAT 101 OF THE COMPILING ABILITY OF MANGANOUS IONS WITH RO3 , CIO^, SO*, HgPO^, F - IONS AND THE DI*SODIUM SALT OF ETHYLSNEDIAMINETETRAACETIC ACID ...... 68

Discussion of the Results Obtained fr«m the Study of the Complexing Ability of the Manganous Ion , 6 9

Separation of Manganese Conqilezes Using Domex 1-X8 R esin 80

Discussion of Results Obtained from the Resin Runs 80

SraCTR0PH0T05ŒTRIC STUDIES OF THE SYSTEM M n'*^-lfaOj ...... 93

Discussion of Results Obtained from the Speotrophotometric Studies * ...... 94

8U1C4&RY ...... 117

Conclusions ...... 123

AUTOBIOGRAPHY...... 125

iT LIST OF TABLES TAB1Æ PAGE

1. Reoorvry of KKnO^ in H^PO^ and HCl Solutions •••••••••• 32

2. RoooTory of KMoO^ in HCIO^ S olutions ...... 33

3* RsooTory of EMnO^ in HgSO^ Solutions ...... 35

4» RoooTsry of lisngsnous Ions ...... 39

5. Rsoovsry Data on the System )bi^-UnO^ •.••••••••••••••• 40

6 * Isotopie Exohange Data ...... 51

7* Increase in Solubility of MnC 2 0 , In the Presence of Various Salts ...... «...... 77

S# In c re a se in S o lu b ility o f RnCgO^ in More H ighly Concentrated Salt Solutions •«••■•••••••«•••••••••••••• 78 ■ 9» Data from Separation of Kn and MoF^ Using Domex 1-16 R e sin 85

10. D ata on L a b ility of M n^ and KnP^ ...... 87

11. Composition of Solutions for Experiments I through XL • 99

12. Optical Density of KMnO^ in Water and Various Acids ... 100

13. Optical Density of DbO, as a Function of Time in the System Mn -IfuOj-Aoid ...... 101

14. Optical Density of EMnO, as a Function of Time in the System ...7 ...... I l l

15* optical Density of KUnO]^ in Various Salt Solutions .... 112

16. Optical Density of KMoOj In Various Salt Solutions •••. 113 LIST OF FIGURES FIGURE PiGE

1. Gugn* Spectrum of Iro n S t r i p Bomb&rded w ith Douteront ...... 22

2. Gmtmm Spectrum of Mangenese Itotopic# Sepereted from Iron Bombarded with 12 Mev* Douteront 22

3* Concentration of Manganete in the Hanganout Ion Fraction After Precipitation with Sodium Hydroxide ...... A4

4# Deoompotition of PotatelAm Permanganate in the Pretence of Vanganout Sulfate ...... 55

5-7« Plot# of ln(l - F) Terse# Tims for Determination of t& • 56

8-10. Plot# of ln(l -F) vertes Tima for Determination of ti •• 57

11-13» Plots of ln(l- F) verses Time for Detormlnation of t^ •• 5Ô

14» Inoreate in Solubility of IfoC^Oj in the Presence of Various Salts ...... 71

15» Inoreate in Solubility of MnCgO, in the Pretence of More Highly Concentrated Salt Sdlttione ...... « 76

16* Inoreate in Concentration of IbCgO Above That in Pure Water ...... 76

17» Comparison of Stability Constants of Complexes of Metal Ions ...... 79

16. Peaks Obtained in the Separation of Different Oxidation States of Manganese Using Dowex 1*26 R esin ...... 91

19» Peaks Obtained in the Separation of Different Oxidation States of Manganese ...... 92

20, Comparison Curves for the System Un'^'^-MDO^-Aoid 114

21. Salt Effects on the Decomposition of ObO^ *.. ». .., ... 115

22. Variation of the Order of Reaction with Concentration .. 116 INTRODUCTION

Appliotlone of Radio*etIvlty In Ch#mietry

The use of redioeotire et oms em treoers in the study of e wide

veriety of ohemloel problems he# beoome an important research technique.

For example, the tracer method is used in the study of isotopie exchange

reactions, structural chemistry, self-diffusion, chemical kinetics, and

analytical chemistry. The high success of the tracer smthod is due

to the fact that in all but the li^test elements, the radioactive

isotopes behave, within limits of the experimental error, the same as the non-radioactive isotopes in chemical and pt^sical prooesses. Also the highly efficient radiation detectors available make possible the estimation of very small amounts of radioactive material, as low as

10"^^ grams in some instances*^ Using multi-channel analysers, the

various components of mixtures of radioactive isotopes often can be identified*

The Problem and I ts H istory

The main problem under consideration is the isotopic exchange reaction between manganous ion and permanganate ion. A literature survey revealed only three researches which dealt with the above reaction. In each the experimenter used radioactive manganous ion as the tracer.

The Mn**-KnO^ exchange reaction was first studied by U. J, Polissar,^ in 1956. The results of his experiments show negligible exchange between

^ah l and Bonner, Radioactivity Applied to Chemistry, New York, John IflLley & Sons, Inc. (1961).

^Polissar, U. J*, J. Am. Cham. Soo. 68 , 1572 (1956).

1 'Un** and UnO^ in 15 minutes at room temperature. The solutions used

In the experiments contained rarlous amounts of low concent ret ions of potassium permanganate (usually about 0 . 02M). perchloric sold and manganous sulfate. The reaction mas stopped by adding sodium hydroxide to the solution mhlch rapidly precipitated most or all of the manganese which was present in the intermediate oxidation states. Polissar also worked on other exchange reactions, for example, manganic oxalate com­ plex ion and UnO^ 1 on exchange and the exchange between the oxalate complex i on with Mn** ion. (Other isotopie exchange reactions were studied in addition to the above but only the Mn -MnO^ exchange will be discussed h e re .)

It sit ou Id be noted here that the addition of smnganous ions to a solution containing permanganate ions produces a precipitate composed mainly of manganese dioxide* The time of appearance of this precipitate in the solution depends on the acid concentration. For a given oon- centration of manganous Ions and permanganate ions, an increase in the acid concentration prolongs the appearance of the precipitate. In neutral solution the precipitate appears almost immediately.

Since Polissar found no exchange between Mn* and MnO^, he con­ cluded that the equilibrium indicated by the following equation does not e x is tf

2MnO^ 4 5Mn** 4 le H* - 5 Mn** 4 8 HgO (rapid equilibrium)

However, it would seem that the following equilibria could exist and still be consistent with the observed resultst ( l) Un04 ^ (rapid rareraible) or

( l l ) 2IÜ1O4 + (rapid rereralbl®)

For this to be true the aasumption must be made that no equilibrium exists between the ions on the right side of equation (I) or equation

(II).

In 1961, A. Tf. Adamson^ studied the Mn**-KnO^ exchange by using a procedure similar to that of Polissar. In Adamson's experiments, dilute solutions of manganous sulfate and potassium permanganate were used. Perohlorlo acid was added to make the final solution approxlsmte-

ly 3M in the acid. By using dilute solutions and excess potassium permanganate the appearanoe of a precipitate (presumably mainly manganese dioxide as suggested above) could apparently be delayed several hours longer than usual. The p erch lo ric sold was used to minindse complex form ation.

Adamson termi rated the exchange by adding sodium hydroxide in excess, then separated the precipitate by centrifugation. He found about Ç>,2% exchange in one hour. From a Vdnetic study of the reaction, a mechanism was proposed idiich involves a rapid pre-equilibrium of the following type I

(III) 5Mn(lI) + Mn(VlT ) - SMn(lTl) + Mn(lV)

The rate-determining step is assumed to be a slow electron exchange between the hydrolyzed forms of the products of the above equilibrium.

The rate equation derived by Adamson is as follows*

R - k(H)* (Mn**)^ (MnO“)®

^Adamson, A. W., J. Phys. Colloid Chem,, 65, 295 (1961). 4 where a, b, and o have the approxlnate values 4/5, 4/s and l/s . The value of k In the equation is given ast

k " 2.6 llt^ nole”^ hr.

As mentioned above, Adamson worked with ezcese permangante in solution. He also tried using excess manganous ion in solution and found complete exchange, but the results were not reliable because of the difficulty of a clean cut separation and analysis. The appearanoe of manganese dioxide apparently stops the exohange completely.

Adamson presupposes that a rapid pre-equilibrium produces two

Intermediate oxidation states (hydrolyzed forms of the products in equation 111 on preceding page) which then undergo a slow exohange.

J. A. Happe and £. S. Martin, Jr.* observed that if the concentra­ tio n of these interm ediates was large, then they must be removed in order to establish a true Mn**-MnO^ exchange. Precipitation 1 s no indication of complete removal of the intermediate substances. Happe and Martin showed that the eonoentration of intermediates had to be large end th e re fo re , unless removed by the ad d itio n of sodium hydroxide

(in Adamson's experiments), could account for the exchange rate observed by Adamson.

In the experiments carried out by Happe and Martin, the exchange between Mn** and MnO“ was terminated by precipitating the permanganate ion as tetraphenylarsonium permanganate. The solutions were approximate­ ly 1-2 molar in nitric acid. Perchloric acid could not be used in the experiments beoause tetraphenylarsonlurn perohlorate is an insoluble salt.

When using nitric acid in the solutions, the precipitate did not appear for about 10 hours. Analysis of the permanganate salt revealed so

^Happe, J ., and Martin, D., J. Am. Chem. Soc. 77, 4212 (1955). 6 activity which corresponded to an exchange half-time of 10-15 hours.

This is not the half-time obtained from the plot of log (1 - P) verses time, but from the plot of £ verses time where E is given as

E - (a/b) (y/x)

Here, a and b are the initial eonoetrtrations of manganous and per­ manganate ions respectivelyi y is the activity of the permanganate fraction at time t, and x is the activity of the manganous ion frac­ tion at the same time t.

The Problem Resulting from Early Work

Speotrophotometric studies of the solutions used by the experi-

Bienters above shoved that before a precipitate appeared, at least 40 percent of the MnO^ was reduced by a reaction approximately first order in MnO^ and w ithout an induction perio d . À so lu tio n 5*6 x 10“^ li in potassium permanganate, 8.7 x 10"^ U in manganous sulfate and 2.0 M in nitric aoio or perchloric acid required about 10 hours for a precipitate to appear. The interim between the tins of mixing and the titoe of appearanoe of a precipitate resulted in a build-up of intermediate oxidation states of manganese in the solution. Their conclusion was that the exchange between Mn and MnO^ was not sufficiently substan­ tiated from the data on the experiments done because the activity which they observed in the permanganate fraotion may not have been due to radioactivity in the MnO^ ions but rather to radioactivity in inter­ mediate oxidation states of manganese that possibly resiain in the permanganate fraction after separation of the two fractions. The main problem of our work was to plan researches that would give definitive 6 data on this problan and also give more infomation aboutthe reactions

InTolved and the species present*

The exchange betsreen Mn** and MnOg has been studied by many experimenters, but there seems to be no consistent agreement in their results* This inoonsietenoy, hom-ver, la of little importance in the

Mn**-MnO^ exchange system, since Adamson reports that the appearance of a precipitate apparently stops the exchange, that is, the exchange between the precipitate and MnO^ is negligible. (Much of the above discussion was taken from the work of Happe and Martin*^)

The MnO^ - IfinO^ exchange has been studied thoroughly by many investigators®*®*^'® and it is generally agreed that a very rapid exohange occurs between these two species* The rate, however, depends on the concentration of base present*

Carrington and Symons® state thatthe rapid electron transfer between MnO^ and MnO^ (and complete exchange) and the rapid interaction 3- between MnO^ and MnO^ in solution siay be taken as evidence for the 3- K reversibility of the MnO^ - MnO^ couple. This could also lead to the conclusion that a rapid and complete exchange exists between 3- ■ ■ MnO^ and MnO^ similar to the MnO^ -MnO^ exchange. This conclusion could eliminate the step Mn(V) —V Mn(Vl) as a rate-determining step.

®Carri ngton and Symons, J. Chem. Soc., 1956, 3373-80*

®Sbeppard and Wahl, J. Am* Chem. Soo., 79, 1020-4 (1957). 7 Hornig, Zimmerman and Libby, J. Am. Chem* Soo., 72, 5808-9 (I960). a Hall and Alexander, J. Am. Chem. Soo*, 62, 3455-62 (1940). 7

Th# rapid electron transfer and the rapid interaction between the species mentioned in the paragraph above indicate that the isotopie

♦ ++ — exchange between Mn and MnO^ most likely occurs by an electron trans­ fer process rather than an atom exohange rocess. The electron change process involves a one (or possibly two) electron transfer between the

++ • various intermediate oxidation states of sianganese in the Mn -MnO^ system. The atom exchange process would Involve the decomposition of 9 a transition state resulting in a rearrangement of chemical bonds.

An exemple is

X - I + X*' * (X— I—X *)' - x" + X—X*

It is often very difficult to determine which process is occurring, but from the above experimental results it is believed that the electron transfer process is occurring in the Mn**-MnO^ system.

Statement of the Problem

The problem at hand, therefore, is to determine whether or not an exchange does occur between and MnO^, If so, the determination of the rate-determining step will also be attempted. A study of some of the parameters that effect the kinetics of the Mn** - MnO^ systems will be made to supplement the above work.

More specifically, the following steps will be taken to determine the rate of exchange and to clarify the processes occurring in the ++ Mn - MnO^ exchangei

1. Determination of the accuracy with vdiioh manganese can be recovered in each of the two fractl' na that are present after separation

®Tiahl 4 Bonner, "Radioactivity Applied to Chemistry," New York, John Mley 4 Sons Inc. (1961). 8 of the Biangaoous i on* from the permanganate Iona, The percent recovery la baaed on a known amount of manganeae dlasolved in the aolutlon com­ pared to the amount determined experim entally. The amount of manganese in each fraotion must be known in order to calculate the apeoifio activities which are used in the rate equations.

2. Determination of the rate of radioactive isotopie, Mn^,

+ +■ — exohange by electron transfer in the system Mn -MnO^ - acid. Two different chemical methods of analysis of the solution will be used.

The f i r s t method involves the use of sodium hydroxide as a sieane of ++ — separating the Mn ions from the MnO^ ions; the second method involves the use of Dowex 1-18 anion exchange resin for separating the two species. Agreement of results obtained from the two different methods will help verify the existence of an isotopie exchange in the system.

S. Determination of the relative stability of complexes formed by Mn*^ Ions in solutions of various anions. The anions to be studied are, CIO^, NO^, 80^, H^PO^, F and the di-sodium salt of ethylene- dlaminetetraaoetic. The extent of complexing with these anions can be used in interpreting the slgnifioance of the + 2 oxidation state of manganese in the rate-determining step.

4. Determination of the oxidation state of manganese in the complex(es) formed in the Mn** -MnO~ -acid systems. Dowex 1-18 resin will be used to separate the various oxidation states. A knowledge of the oxidation states present in the Mn**-linO^ system can help in deciding the rate-determining step. &• Determination of the time of appearance of manganeme dioid.de in the }Ad ** -UnO^ -acid eyitem using a speotrophotometer to follow the deoompoeition of permanganate and to note the time of appearanoe of manganese dioxide in the system. The kinetics of the system will then be studied in order to determine the meohariism inTolwed in the iactopio exohange.

Stability of farieaa Oxidation States of Manganese

Manganese exists in all oxidation states from zero to +7. Many of the lower and nearly all of the higher oxidation states of manganese form highly colored compounds. Preparation of oompounds of each of the seven oxidation states has been reported, but some of these states are extremely unstable. Eaoh state will be discussed in turn starting with the unlpositive oxidation state.

Manganese (l)

Manganese is reported to exist in the +1 oxidation state in the compound Kg Mn(CN)g . It has been prepared eleotrolytically by

Treadwell and Raths.^*^ Others, hare also reported preparing the compound, but certain evidence such as magnetic p ro p erties has cast some doubt on the v a lid ity of the composition of the compound. In any event, the compound is considered a powerful reducing agent and is not stable in solution.

^Treadwell and Raths, Helv. Chim. Acta, 35, 2259 (1962).

^^Manchot and G all, Ber., 60, 191 (1927).

^^Grube and Braueh, Ber., 60, 2275 (1927).

^®Klemm, Angew. Chem., 63, 396 (1951). 10

Manganèse ( il )

The +2 oxidation itate of manganese Is one of the two most stable states of manganese in neutral or slightly aoidio solution. In basic solution, the manganous ions are precipitated as Mn(OH)g. This compound when heated, converts to MnO which is stable under certain experimental conditions.M anganous hydroxide is easily converted to MnD(CE) in the presence of oxygen. In the presence of higher oxidation states of manganese (+4, +5, +6 and +7), the manganous ions act as reducing agents,

Many compounds exist in which manganese is present in the +2 state;

MnClg, MnCgO^, MnSO^, MnClO^, Mn(MOg)g (and hydrates of th ese compounds), and K^Mn(CM)g' 3E20 are just a few. Dissolving the metal or a precipi­ tate of manganese dioxide in hydrochloric acid results in the formation of manganous chloride as one cf the products.

Only two other states of manganese compare with the +2 oxidation state in stability. These are the +4 state as manganese dioxide end the +7 state as permanganate ion.

Manganese ( i l l )

In general, the +3 oxidation stste of manganese .Is very unstable in solution unless in the form of a complex. The +5 state forms stable complexes in solution with chloride and fluoride ions.Th: a state is also stabilised in ooncentrated sulfuric acid solutions but upon dilu- 16 17 tion the +3 state reverts to the +2 state. ' In aqueous solution the reaction Isi

MoAlpine and Soule, "Q u alitativ e Chemical A nalysis," New York D. Van No at rand Company In c. 1955.

^^einland and Dinkelooker, Z, anorg. chem., 00, 173 (1908). ^^French & Howard, Trans. Faraday Soo., 52, 712-16 (1966).

^^Domange, Coiapt. rend., 208, 284-5 (1939). 11 +s *2 + 2 Mn ^ + 2 EgO - UnOg + Mn + 4E

Thus tervalent nanganese Is favored In aoidio solution, Except for

WngOgand UnOOE, te rv a le n t xoanganese compounds are much less stab le than the b iv alen t.T h e compound MnClg exists and is stable In the absence of moisture. In the presence of the latter, the compound 19 décomposés into manganous chloride and chlorine.

Thus, in sunning, it may be said that the existence of the +3 oxidation state of manganese in solution is transient and depends upon the presence of complexing agents for any appreciable concentration to exist. The concentration of manganic ions (Mn^*) in the absence of complexing agents may be assumed quite lorn.

Manganese (IV)

By far the most stable form of the +4 oxidation state of manganese is manganese dioxide. This oxide is a powerful oxidant and is readily reduced to Mn(ll). No manganese(lV) salts have as yet been characterired,^ although compounds such as KgHnFg, KgUnClg and Mn(SO^)g are reported to 18 exist in crystalline form or in very concentrated acid solutions.

However, a ll of the above compounds are quickly hydrolysed to MnO g by water. Therefore, it appears that the existence of Mn(lV) in solution is transient and practically negligible ; the precipitate MnOg always appears to be favored.

1 Q Partington, "General and Inorganic Chemistry," London, UacîÛllan and Co. L t'd , 1968.

19 Chretien and Varge, Bull. soo. chim. 6 , 3, 2385-94 (1936).

20 Moeller, "Inorganic Chendetry," Now York, John Wiley and Sons Inc. 1962. 12

Uang»ne«e(V), (Vl) And (Vll)

Carrington and Synons® hava prepared the +5 and the ‘♦6 s ta te s of manganese and have been able to obtain speotrum on each species*

They have found that each species follows Beer's law in dilute solutions,

They also point out that slnoe the hypomanganate Ions do follow 5- Beer's law, the ions most likely have the formai la HnO^ rather than

MnOg since the latter would have a high tendency to associate in solu­ tion and thus deviate considerably from Beer's law.

A hypomanganate solution is blue in color* In 12 M potassium hydroxide, the ions are stable up to 3 days time and then rapidly decompose. The appearanoe of manganese dioxide greatly accelerates 5- the rate of decomposition. The authors state that the MnO^ ion is remarkably stable as indicated by the feet that hypcatanganate can be prepared simply by heating manganate in concentrated alkaline solution, and that continued heating of the resulting blue solution to high temperatures for long periods of time will not alter the composition of the solution. (Lux and Niedermaler^^ found that high temperature favors However, once conditions of protonation occur, the hypomanganate rapidly dlsproportionates forming manganese dioxide as one of the products. Thus, the concentration of MnO^^ in acid solution may be considered to be transient and essentially negligible.

Manganate solutions are green in color. A solution of potassium manganate in 4M potassium hydroxide Is stable for several

^^Lux and Niedennaier, Z. anorg. u. allgem. chem., 285, 246-61 (1Ô56). IS woeks. Uangaaate like hypomanganate disproportionates in aoid solution.

Therefore, the eonoentration of stanganate in aoid solution may be

oonsidered to be low* Manganate is easily prepared by adding sodium

or potassium hydroxide to a solution of potassium permanganate until

strongly basic and then heating if necessary.

Permanganate (MnO^) solutions are purple in color. Potassium

permanganate in slightly acidic solution is stable for several months if

kept in the dark. Permanganate can be prepared by oxidising manganous ions with a strong oxidising agent such as IgO or KIO 4 , The la tte r oompound requires heating in solution to oxidise the manganous 1 one to permanganate ions. The former does not.

The +7 state of manganese as permanganate ion and the +2 state as manganous ion are the two most common and stable states of manganese under ordinary conditions.

Hence the following reactions stay be considered*

+5 ++ -2 +6 -2 +7 _ MnO ^MnO~ ^ MnO ^MnO^ Stable y y Concentrations alone May be j y low in solution; existing fleeting most MnO , slow stab le of th ese f le e tin g ones e'* values for these reactions are as follows (in slightly aoid solution) :

1.18„ 2+ -1.5 „ 5+ -0.95 Mn Mn Mn' MnOglÊlîL MnO^* ibiO" -1.25 -1.70 14

also

Mn(CN)g" Mn(CN)^~

It i« »e«n from the E° Talvies that the higher oxidation states of manganese ebo-re the *2 state are not favored in dilute aoid solution*

However, if oomplexlng anions are present (e,g. CN , P , EDTA etc.) the electrode potentials are changed and in fact, in the case of CN ion

complexes of the +2 and +5 states, the potential becomes positive indicating a favored +5 state. No data could be found for the electrode potential involving complexes of the +4 state of manganese* Experiments

carried out in this research indicate that the complex involving the

+5 state of manganese is highly favored in the system -MnO^ -EF, that is, the following reaction predominate si

4 Mn*'*’ + MnO" + 30F" + 8 E* « 5 MnP®* +4E 0 4 6 2

4* 2» If this is true then it appears that the couple MnF. —MnF has a c 6 o 4" 5* loner E value than the couple MnFg * EXPERIMENTAL PROCEDURES

Preparation >ad Purification of R#&g#nt#

Ail ohemlcals uged in the study of this problem were of Reagent

grade. No further purification was attempted. Double-distilled

demineralized water was used in all solutions. The main solutions

used in the experiments were prepared as indicated below:

Potassium Permanganate Solution

Potassium permanganate was used in a ll of the Mn^^-MnO~ exchange

experiments and in the solutions used in the speotrophotometric studies,

The main s o lu tio n of potassium permanganate was made up as follows *

A known amount of potassium permanganate was placed in a 1-liter

volumetrlo f la s k and d ilu te d to the mark. The so lu tio n was mixed

thoroughly and transferred to a larger flask. The solution was boiled for about 15 minutes to insure oxidation of any impurities in the

solution. It was allowed to set for approximately two days and then

filtered through an asbestos filter on top of some glass wool placed in a small funnel. The solution was placed in an amber colored bottle

and stored in the dark when not in use.

îianganous S u lfa te Solution

The +2 oxidation state of manganese used in the Mn*^-MnO system 4 was in th e form of manganous su lfa te . The follow ing prooedure was used in preparing this solution:

A small crucible was placed inside a larger crucible lined with asbestos fibres. The crucibles were heated for about 50 minutes with

16 16

the mangaaouB su lfa te la the Inner oruolble at a tem perature of 450 to

500° C. Care was taken to avoid over-heating the salt as this will

cause the manganous sulfate to décomposé into a rust colored powder

whioh appears to be insoluble in water.

A weighed quantity of the above salt was plaoed in a 1- l i t e r

flask and diluted to the mark. After setting several days, a very

thin layer of the rust colored powder appeared on the bottom of the

flask* The powder was believed to be an oxidized s ta te of manganese,

possibly MnOOH. The solution was filtered through filter paper and

plaoed in another flask. A small correction was applied to the

solution of manganous sulfate beoause of this loss. Molarity of this

solution was 0.0492M. Another solution of owinganous sulfate was also

made by f i r s t heating the s a lt at 440-460° C. for one hour; no

precipitate appeared in the solution propared from this sample. The

caloulated molarity of the latter solution (0.0501'4) agreed quite well

with the observed value (0.05031).

It should be mentioned here tlmt two potassium permanganate solu­

tions were prepared for these experiments. They v/ere frequently checked

and cross oheoked using the sodium oxalate method and the iodine method.

The latter meti od is described be low. The molarities of all important

solutions used in the experiments were oheoked with calculated values sdie never possible and some re-oheoked to insure proper values of molarities. The solutions used and their ooncentrations will be indicated as the experiments are discussed. 17

Prooedure Uaed fo r Reooyery of Mapgaaou» Ion»

The eonoentration of nanganoua ions in solution was determined by oxidising the «manganous Iona to permanganate Iona using silver oxide (AgO). The eonoentration of permanganate was then determined by the iodine method deaorlbed later.

The so lu tio n oootaining the manganous ions should have no ohlorlde, bromide or iodide ions present. If these lone are present in solution, they will form a precipitate with the Ag* ions which appear from the siIver oxide treatment. The formation of the precipitate hinders the oxidation of the manganous ions. Therefore, the halogen lone should be removed by evaporation with aulfurio acid. In the absence of these

Iona, th e following prooedure was uaedi

The solution was made approximately IM in sulfuric aoid, then cooled to ro'jm temperature or lower. Small amounts of silver oxide were added slowly, with frequent swirling of the flask, until a small excess of the oxide had been added to the solution. A large excess was avoided beoause of the iodine treatment which follows. However, th e amount of excess s ilv e r oxide appeared not to be c r i ti c a l as long as a moderate amount was used. After the silver oxide treatment, the solution was made 5M in sulfuric aoid. The solution was slowly heated to 70-90° C. or near the boiling point but was not boiled. After setting for at least 15 minutes, the solution was cooled to room temperature or lower. Then iodide ion was added to determine the con- oentration of permanganate ion present. The procedure la described below. 18

Prooedure for Detoradnlag Pera*nganate Ualag Iodide loa

The permenganate solutloa muet be made aoidio if not already ao.

The aoid used should be sulfuric aoid or an acid which will not decom­ pose in the presence of permanganate over a period of several sdiiutea.

The solutions used in the reoovery prooedure here were about 1-5 molar in aoid. Hydroohlorio aoid oan be used if the aoid is kept dilute and the titration is made immediately* After adding the aoid and

oooling, if neoessary, potassium iodide salt is added until an excess is present. If the Silver oxide treatment above Is used, more potassium iodide must be added to precipitate the Ag* ions as silver iodide. The iodine liberated is titrated with sodium thiosulfate using starch

solution as an indiontor. The end-point is very sharp in the absence of silver iodide and quite reproducible. In the presence of the precipitate the end-point is not so sharp but still detectable.

Preparation of Radioactive Manganese, Mn^^

Mn^ has a half-life of 300 days. It is a gamma emitter. It can be obtained by bombarding Fe^® with douterons. The reaction is Fe^^

(d, ) Mn^. Other radioactive elements are also formed in bombarding

Fe^® and must be removed from the Mn^ since pure manganese compounds are used in all experiments.

Iron Target Bombarded with Douterons KMev.

The following reactions may ooourt

= U n^ + gHe^ Main rea ctio n

t i / 3 " 300 days 19

Fe56 + _ p*57 +

= Fa®® + abundaaoa about 6ji only

Fa®® + — Co®® + 2n high p ro b a b ility

Fa®^ + = Un®^ + gHa*

t i / 2 " 5.6 days

These are the main reactions that take plaoa in the iron target. Other réactions may take place but they are of little Importance.

The following prooedura was devised to obtain Mn®^ from an iron ta r g e t t

A piece of high purity iron strip was bombarded with douterons at the Ohio State University cyclotron. The time of bombardment was about 3 hours. The aotivated iron strip was tiien placed in a concen­ trated solution of hydroohlorio acid for about 3 hours. The iron strip was removed, washed with double-distilled water and checked on the monitor for activity. The monitor indicated about 20,000 counts per Tainute at the surfttoe of the tube. This is only a small fraction of the aotirity originally present in the iron. Thus, by dissolving the outer layers of the iron strip, most of the activity was removed.

To the acid solution containing most of the activity was added

2.0 00 of O.lOU manganous chloride and 2.0 co of O.lON oobaltous chloride and the resulting solution then stirred and conoentrated 20

hydr^t:hlorlo aoid added until a dark green solution appeared. Con­

oentrated hydroohlorio aoid was added in exoess*

The solution was than poured into a small separatory funnel. The

stem of the funnel was attaohed to an anion e^ohange oolumn containing

Dowex 1-X8 resin (used in all resin experiments in this research).

The resin used for the above separation was first treated with con­

centrated hydroohlorio aoid before the aotive solution was allowed to

drip onto the column. Previous experiments, with concentrated and more

d ilu te hydroohlorio aoid so lu tio n s, on the re sin showed th a t concen­

trated t^rdroohlorio aoid causes the solution from the oolumn to be a

pale yellow. This, however, did not affect the separation. In 6N

hÿ'droohlorio aoid the solution from the oolumn was clear.

The green color in the aotive solution remained at the top of the

oolumn. This band contained the oobalt ions and ferrio Ions slnoe

these ions form strong complexes in the presenoe of 6 N and more con­

centrated hydroohlorio aoid solutions. This conclusion can be drawn

from the apeotrum discussed later. The eluent from the oolumn, however,

was not olear but very yellow, indicating ferrous chloride ooming through. The Mn*^ and Fe** ions do not form strong enough complexes

in 6N hydroohlorio aoid to remain on the oolumn and thus are elu ted when the column is washed with 6 N aoid. The column was washed three times with 6 N hydroohlorio aoid. Periodic oheoke were made until

activity oeased to come from the oolumn. The green band moved very

little during this washing process. It was then removed with 0.5N hydroohlorio aoid and the resin checked for activity. None was deteotable. 21

The yellow solution above was tw o treated with SgOg to ozidiee the Fe** to Fe***. The solution was oarefully evaporated on a hot plate while blowing air on the surface of the liquid. At a volume of about 50 ml, the flask was removed, the solution oooled and again poured through the ion ezohange oolumn. Previous to this, the oolumn had been treated with oonoentrated hydroohlorio aoid again. The yellow band remained at the top of the oolumn. This was the ferrio chloride oomplex. The solution coming from the resin was olear. The very pale yellow color present may have been due to slight decomposition of the concentrated hydroohlorio aoid in the column as mentioned above.

After the solution had passed through the column and the latter washed with 6N hydroohlorio aoid, a check was made on the activity of the yellow band at the top and it was found to oontain praotioally no activity compared to the fraction already removed. This is as expected since the additional reaction in the target, Fe^® (d,p) Fe®*^, involves only stable isotopes. The other target reaotlon, Fe^* (d,p) Fe^^, is not a very probable reaction (the abundance of Fe®^ is about 6^ ).

The yellow band was removed with 0.5N hydrochloric aoid. By use of a portable Geiger counter it was observed that the ions in the green band contained much more of the total activity than the solution washed through the oolumn, although the latter was also quite aotive

(it contained the Un^) . This, too, is to be eipeoted since the probability of Fe^^ (d, 2n) Co^ is much greater than Fe^® (d,«() Mn^ reao tlo n . 22

A g&mmm my •peotrus iras obtained on the Iron strip before it araa dissolved, and on the solution after the iron surface had been dissolved.

The peaks ooourred at the expected places for a mixture of oobalt and manganese isotopes. Then a garnna ray spectrum mas run on the solution obtained frcn the oolusm ahioh supposedly contained only the manganese isotopes. The speotrum showed peaks at 0.49-0.51 Mev (this is annihila­ tion radiation from f*), 0.74 Mev, 0.82 Mev, 0.93 Mev, and 1.40 Mev.

This prominent peak at 1,24 Mev from Co was absent. This peak did ooour previously in the spectrum from the original solution before passing it through the ion exchange oolumn. Figures 1 and 2 show the two sp ectra before and a f te r separation for comparison.

The counter %as standardized using known gamsia standards. The absence of this prominent oobalt peak would indicate a good separation of oobalt from the manganese fraction. All of the peaks in the manganese gamma ray spectrum can be accounted for if the isotopes

Mn^^ with a 300 day half-life and with a 5.6 day half-iife are assumed to be present. The Mn®^ comes from the reaction Fe^ (d,fl( }

and the from F e ^ (d,o( ) Un^^. A check on the modes of decay of these two isotopes shows gamma rays emitted at the energies men­ tioned above. A much later spectrum of the sample showed only the

0*84 Mev peak from Mn^. The 5.6 day Mn^^ had decayed.

Yield of Mn^ from the Reaction* Fe®® (d,4) Un®^

The crystal used to determine the yield was a 2-l/2 inch sodium iodide crystal with a 5/s by l-l/2 inch well in the center. The standard was a Mn^ gaiana source which gave 1 .6 x 10^ j[/ain. The discriminators 224 lOO

•o

•0 TO

• 0 oI 40 K>

to

030 040 I to 130 110 f TO 330 940

Ftgur* I. Gommo-roy ipoctrum oMoinad from bocobo'dmarrt of high punty moo with dautaront

lOO

*o #0

TO

«0

I ^

40

90

to

0 090 040040 130

Figura 2 Gommo-roy sfMCtrum of MongonaM itotopa* taparofad from daufaron bombordmant of iron. 23

oa the counter were met at the lower edge of the Mn^ peak. At thlm

setting the standard source gave 1,950 CPH. At the same setting, 1 ml.

of the manganese separated from the ion exchange oolumn gave 16,460 CPH.

Assuming the geometry was the same in both samples, then 1 ml. of the

sample had 1.56 X 10® i^nln. There was a total of 26 ml. so total

yield was 3.54 X 10^ ^adn or 5.9 X 10® ^seo. There are 5.7 X 10^

disintegrations per second per ndorocurie. The yield was therefore

15.9 mioroouries. Efficiency of crystal used was calculated to be 7.9#.

Experimental Procedure for Studying the Isotopio Exchange Reaction

Between Mn*^ and MnO*

The solutions of potassium permanganate and manganous sulfate

were made up to the desired concentrations and the flask placed in a o constant temperature bath at 25 C. The ti e was noted when the

potassium permanganate was added to the solution. At certain time

intervals, one m illiliter of the solution was removed and made

slightly basic with sodium hydroxide. A precipitate formed immediately.

The mixture was centrifuged for 3-5 minutes and the liquid withdrawn

with a very small tipped eye-dropper. In most of the early runs, the

liquid was passed through a fritted disc filter to insure complete

recovery of the precipitate. This liquid, which contained the

potassium permanganate, was made acidic and the permanganate was

determined by the Iodine msAhod. These solutions were partially

evaporated for counting purposes. The precipitated fraction was not

checked for recovery, but non-radloaotive runs were made later under identical experimental conditions and both fractions were checked for 24

rsooTory* The precipitate was diaiolTad in conoentrated hydroohlorio

aoid and diluted with water. Since the to I udos of the Mn*^ fraction

and MnO^ fraction were about the lamo, a correction was not necessary

in the counting rates.

YThen the anion exchange resin was used in separating the two

fractions, the same procedure was used except that the reoowary of both fractions was attesçted by using the silver oxide method.

Experiswntal Procedure to Determine the Time of Formation of Manganese

Dioxide for Studying Kinetics of the Mn*^-MnO^ System

Early experiments showed that a disturbing factor in the study

of this system is the formation of a precipitate of manganese dioxide,

MnÛ2 after the exchange reaction has been in process for a short time.

The formation of the solid appears to stop the exchange and this makes it impossible to follow the reduction of the permanganate spectro- photometrioally. Consequently it was desirable to know the rate of formation of the precipitated MnOg under different experimental conditions.

Since the Beokmann Model DU Spectrophotometer was used in this study, the solutions were made up simply of varying concentrations of all species present, that is; Un , UnO^, and whatever aoid was being used. A reading was taken at certain time intervals on the spectro­ photometer . The rate of decomposition of permanganate was studied in the presence of five different acids; namely, perchloric acid, nitric aoid, hydrochloric acid, sulfuric aoid, and phosphoric aoid. The rate was also studied in the presence of the salts of these acids at 26

approximately the same pH aa the original aoid solutioae. The pH

was determined with a glaee eleo tro d e. 5- ■ The appearanoe of colored epeoles such as UnO^ and HnO^

should not effect the optical density readings much since they are in

such low concentrations, especially in add solution.

Experimental Procedure to Determine Iftfhether Manganous Iona Form

Complexes

Manganese in the +3 oxidation s ta te forms complex ions re a d ily but not much information is available on the relative oomplexing tendency of manganese in the +2 state with inorganic anions, other than

it Is presumed to be much less. The following experiments were devised to determine the oomplexiag tendency of Mn** ions.

Manganous oxalate was used in these experiments. The solubility is approximately 0.3 gram per l i t e r . The s a lt was obtained by mixing manganous chloride with oxalic aoid (or sodium oxalate) and allowing the solution to set for about one week. By the end of this time well defined crystals of sianganous oxalate (hydrate) had appeared. The salt was washed thoroughly several times and small amounts were placed in five different flasks. A measured sunount of water was added and the flasks were put in a constant tempeirature bath at 24.0 + 0.5° C. and stirred. Samples were withdrawn after the flasks had been in the bath from 24-48 hours. Standard salt solutions were added to the water in the flasks. From the data and the amount of water in the flask and the quantity of salt added, the concentrations of salt in the solutions was calculated. The effect of the five salts on the solubility of manganese 26

oxalate «as determined* sodium perohlorate, sodium aitrate, sodium

fluoride, sodium di-hydrogen phosphate, and the di-sodium salt of the

EDTA oomplexing agent. After ssnspling as described above, more salt

solution was added to the proper flasks. The solubility of the manganous

oxalate was determined in four or more oonoentratlone of each of these

s a lts .

Since oxalate ion Is a fairly good oomplexing anion, the data

obtained from these studies was purely relative to the oxalate oom­

plexiag ability.

Preparation of Anion Exchange Columns

Anion exchange oolumns were used to separate various oxidation

states of manganese in the Mn**-MnO^ system. The solutions were passed

slowly through the column. Small samples were collected and counted

for activity present. The columns were prepared as follows*

The column was first filled with water to about two inches from the top. Any bubbles present in the tube were carefully removed.

Glass wool was packed in the bottom to prevent the Dowex 1-X8 resin from passing through. The resin was 100-200 mesh. TShen u sing hydro­

fluoric aoid solutions, a plastic tube was used for a column and cotton was packed in the bottom since the hydrofluoric aoid would dissolve the glass wool in a few m inutes. The cotton seemed to be unaffected by the hydrofluoric aoid solutions.

A thick paste-like mixture of the resin was made up and slowly poured in to the top of th e oolumn and th e oolumn drained slowly to aocononcdate th e added volume. The resin was added to about 2-3 inches 27 from the top of the column* The reein was then baok-treshed for sereral minutes to insure an even paoldng of the column and to eliminate any air bubbles that may have entered while adding the resin* Glass wool was sometimes placed on top of the resin to prevent stirring up the re s in when solutions were being added*

Dowex 1-X8 resin expands and contracts slightly with various con­ centrations of aoid. This disturbance apparently does not affect the ion separating ability of the resin as long as excess solution is kept on top of the resin. THEORY

The equations used In this problem and the assumptions used in

their derivation, if any, are as follows*

Consider the following general type of exchangej where A* and B*

represent radioisotopes of the same element in different oxidation states

and A and B are the stable speoles;

A* + B = A + B*

The important equation used in this type of isotepic exchange

reaoticn is the following!

where (A*) is the concentration of radioactive X (X is the element

manganese) in the A form at time t, (B*) the concentration of radioactive

X in the B form at time t, (A) and (B) are the total concentrations of

X in the A and B form respectively and R is the constant rate of ex­

change of X atoms between A and B. Integration of the equation gives

fe (1 - P ) . - where F for a homogeneous system is equal to

(A') - (A'J (S') - (B* ) F (A’oo)-U’o) (B'oo)-(B'o)

The subscripts refer to time zero and infinity.

The important assumption is that the concentrations (A) and (B) do not change appreciably during the tine of the experiment. If this

28 29

is tri)e, the integrated equation la applioahle to the exchange data

obtained from the expérimenta. Correctiona applied to the F factor

arill be diaouaaed in the treatment of data aeotion.

In the atudy of the decompoaition of permanganate, the equation

uaed i a

d ■ kcl

In thia equation, d la the optical deneity, k ia a constant, c is the

conoentration of the solution under study and 1 is the path length in

centimeters through the solution* Written in differential and the

logarithmic form the equation la

- - kcl

and j log T“ = - kcl 0

1 is the intensity of light at path length 1 in a solution of concen- tation c, is the intensity of light at the point of entrenoe into the solution.

From the optical density measurements, the conoentreti on of per­ manganate at various times can be calculated. From these data, it may be possible to derive a rate equation of the folloning type for the décomposition of permanganate in the presence of manganous ions

j_ - a ++ b - ^ - k’ (Kn 04) (Mn )

The reaction depends on the hydrogen ion concentration. Therefore, if the hydrogen ion is not in large excess, k' is equal to k(H*). 30

In the «tudy of the oomplexing a b ility of the amnganoue ion, the

Debye - Eüokel equation wee applied to the data obtained fron very dilute Bolutiona. The equation is

log f * -A £+1 \JiX

This ie the limiting equation of Debye and Huckel whioh ie applicable only to very dilute solutions* EICPERIMEMAL DATA AND RESULTS ON RECOVERY OF MANGANESE

Molarity of Solution* ü«ed In Eaperlwfnts

The molarities «111 be given as the experietents are dlscuBsed.

None of the solution* appeared to have deoompoeed, but evaporation caused a slight change in molarity In some of the solutions over a period of several months. This resulting change in ccncentrati on was determined by restsndardising the solution.

Recovery Data on Permanganate Solutions

In all of the samples in which a determination of the concentra­ tion of manganese was made, the general procedure aas to oxidise the manganese to the +7 oxidation state, if not already in that state, and determine the concentration of permanganate by the iodine method (as described earlier). The term "recovery," used frequently throughout this thesis, is used to indicate a comparison between the experimentally determined concentration of manganese in a sample end the calculated or true concentration of manganese (in the same sample) as determined from standard values obtained in the preparation of the solutions.

Recovery of permanganate in water, in different acids, and in different concentrations of these acids was thoroughly explored by the iodine method. Tables, 1, 2, and 5 give the results of this work.

In every sample in each of the three tables, approximately 1.0 ml. of UnO^ was used. A sample of 0.975 ml. was used in samples 10-22 in Table 1, and 1 through 10 in Table 2. The "Observed Molarity" values given in Tables 1 and 2 for these samples are corrected to

31 &2

T able 1

Recovery of Potaoslum Permanganate in H 3PO4 and HCl Solutions

Sample Experimental Ml. Observed* , Calculated Number Conditions NagSgOg M olarity x 10 Molarity x 10' KMnO^ KMnO*

1 . Original 3.IN HgPO* Solution 2.05 4.10 4,10 2 . Same as sançle No. 1 2 .02 4.00 4.10 3. Conoentrated HgPO^ Solution 1.96 3.92 4.10

4. Original 3.19 HgPC^Solution no water added 2.08 4.16 4.10 5. Same as No. 4 plus 1 00 water 2.04 4.08 4.10 6 . Same as Sample No. 5 2.04 4.08 4.10 7. Same as Sample No. 5 plus large amount water 2.04 4.08 4.10 8 . Same as Sangle No. 5 plus 1 drop 0.05M MnSO^ ( a f te r 20 m ln.) 2.04 4.08 4.10 9. Left over 3.IN HgPQ^solution 2.06 4.12 4.10

10. Original solution (oootained no aoid) plus 1 drop oono. HCl 2 .0 0 4.10 4.10 11. Original solution plus 1 drop oono. HCl plus 4 00 w ater 1.99 4.08 4.10 12. Original solution plus 4 drops oono, HCl. no water added 1.55 5.18 4.10 13. Same as No. 12 except 8 drops oono. HCl 1.94 3.98 4.10 14. Same as No. 12 one hour later 0.63 15. Left over KMnO^ solution 2.14 4.59 4.10

16. Original solution (contains no aoid) plus 1 drop oono. HCl 1.99 4.08 4.10 17. Same as No. 16 plus 4 00 water 2 .0 0 4.10 4.10 18. Same as No. 17 plus 8 drops oono. HCl 1.53 2.73 4.10 19. Same as No. 17 plus 1 drop 0.05H MNSO4 2 .0 0 4.10 4.10 2 0 . Original solution plus 1 drop oono. HCl. 2 drops MnSO* 0.06M ( a f te r SO mln.) 1.96 4.02 4.10 2 1, Original solution (contains no aoid) plus 2 drops oono. HCl 1.97 4.04 4.10 2 2. Same as No. 21 1.96 4.02 4.10 *Conoentrations in all Tables in this research given In moles per liter unless stated otherwise. 33

T able 2

Re0ovary of Potaselun Permanganate in HCIO^ Solution*

Sangle Expérimental Ml. Obser. I lb 04 Calo. KMJ1O4 Number Condition# ^*2^2^ Cono. X: 10® Como. X 10®

1. Original 3.IN HCIO4 to lu tlo n plus mater 1.90 5.90 3.94 2 . Same a# No. 1 1.95 3.96 3.94 3. Same a# No. 1 but no mater 1.91 5.92 3.94 4. Cono. HCIO^ solution 1.76 3.61 3.94

Ô. Original 3.IN HCIO^ solution plus 0,25 ml. mater 1.93 3.96 5.94 6 . Seme a# No. 6 but 1 ml. mater 1.97 4.04 3.94 7. Same a# No. 5 but 0.5 ml. mater 1.92 3.94 3.94 8 . Same as No. 5 but large amount of mater 1.90 5.90 3.94 9. Same as No. 5 but 1 drop 0.05 M MnSO^ (30 min. later) 1.80 5.70 3.94 10. Left-over solution 2.05 4,20 3.94

11. Original 3.IN HCIO^ solution plus 4 drop* oono. HgBO* 2.07 4.14 4.16 12. Same a s No. 11 but no BgSO^ 2 .1 0 4.30 4.16 13. Same a s No. 12 2 .1 0 4*20 4.16 14. Same as No. 11 2.08 4.16 4.16 16. Same as No. 12 plus mater 2.08 4.16 4.16 16. Left-over solution plu* mater 2.09 4.18 4.16

17. Original 5.IN HCIO^ solution plus 3 ml. mater 2.04 4.03 4.10 18. Same as No. 17 plus 6 drop* oono. NgSO^ 2.04 4.08 4.10 19. IgO treatment on original solu­ tion plus 5 drops oono. HgSO^ and mater 2 .04 2 0 . Same a s No. 19 2.05 4.10 4.10 2 1 . Same as No. 19 2 .0 2 4.04 4.10 2 2. Same as No. 17 2 .0 2 4.04 4.10 23. Left-over solution 0 .2 0

24. Same as No. 19 2.06 4.10 4.10 25. Same a s No. 19 plu* 6 drops oono. HgSO* 2.06 4.12 4.10 26. Same as No. 25 2.06 4.12 4.10 27. Grig, sol'n * 0.285 ml. 11.IN NaOH. filtered, washed. 8 drops oono. HgSO^ 1.94 28. Same as No. 27 but not filtered 1.95 3.90 4.10 29. Left-over solution 2.05 4.10 4.10 34

Tabla 2 Contd.

Sample Expérimental Ml. Obser. KMnOa Calo. EHnOg Number Conditions Na 2®2°5 Cono. X 10* Cono. X 10*

1 . a Original 3.IN BCIO^ solution 1.98 3.96 3.96 2. Same as No. 1 1.98 8.96 8.96 3. Same aa No. 1 1.98 3.96 3.96 4. Same as No. 1 1.98 3.96 5.96

5, O rig in al 5 . IN EC 10^ so lu tio n , AgO treatment, heated, oooled 15 sdn. in air than to room tastp. in water, 2 drops 4^ NaCl solution added 2.06 4.12 4.10 6 , Same as No. 5 except no NaCl so lu tlo a 2.03 4.06 4.10 7, Same as No. 5 1.96 5.96 4.10 8 . Original solution plus 5 drops water & oono. EgSO^ 2.05 4.06 4.10 9, Seme as No. 8 except no aoid 2.03 4.06 4.10 10. Same as No. 8 2.02 4.04 4.10 11. Left-over solution 0.05

i s ! Original 5.IN HCIO^ sol'n plus water 1.95 3.90 3.96 13, Same as No, 12 ten minutes la te r 1.94 3.88 5.96 14. Same as No. 12 20 ml mutes la te r 1.92 5.84 3.96 15. Same as No. 12 30 minutes later 1.91 3.82 3.96 16. Same as No. 12 40 minutes later 1.90 5.80 3.96

• Pura 3,96 z 10"^ M EMnO^ aolutloa usad as blanks for samplas 12 to 16. * la addition to balog S.96 % 10"3 M In KMnO^, samples 12 through 16 also wara 6 ,8 z 10"* M lu MnSO^, No attampt was mada to reoover the manganous ion fraction, only tha panaanganata ion la a 1 .0 ml, sample was raoovared.

An attempt to recover MnO^ ions in perohlorio acid solution siftor

passage through HCR-12 cation exchange resin was also triad, but even at a flow rate of 1 ml. per minute, considerable decomposition occured.

T herefore, th is resin was not used. s e

Table S

Recovery of Poteeelum Permanganate in H 2S0^ Solutions

Sample âcperlmëntaï ÎE% Observed Calo'd KMnO* Number Conditions Cone, % 10^ Cono. % 10^

1. Original S.IN E SO solution only 2.04 4.0# 4.10 2. Sane aa No. 1 2.06 4.12 4.10 5. Sane aa No. 1 2.02 4.04 4.10 4. Conoentrated K^SOg solution 1.72 3.44 4.10

5. Original 3 .IN EgSO* solution only 1.90 5.98 5.94 6 . Original 5.IN EC 10^ solution 1 ml. plus 1 ml. water 1.96 5.90 5.94 7. Original 5.IN EC10 solution 1 ml. 1.97 5.94 5.94 8 . Original sol'n plus large am't water 1.98 5.96 5.94 9. Original solution plus 1 drop 0.05W UnSO. (26 min. la te r) 1.96 5.92 5.94 10. Left-over solution 1.91 5.62

11. Original 5.IN EgSO* solution plus 8 drops oonc. H 2S0^ 1.86 5.76 12. Same as No. 11 except 4 drops acid 1.88 5.76 13. Original sol'n + 0.29 ml. 11.IN NaOH plus 4 drops cone. KgSO^ 1.87 3.74 14. Same as No. 15 except 6 drops aoid 1.92 5.84 15. Same as No. 15 except 9 drops aoid 1.92 5.84 (also added 3 ml. water to Nos. 18 4 19) 16. Original solution plus 6 drops cono. E2SO4 1.91 3.82 17. Left-over solution 0.83

16. Original 3.IN E 2SO4 solution plus 5 drops conc. H 2SO4 2.04 4.08 4.10 19. Same as No. 16 except 6 drops acid 2.06 4.10 4.10 2 0 . Same as No. 19 (45 mln. later) 2.04 4.08 4.10 2 1 . ÀgO treatment of original solution 2.04 4.08 4.10 22. A fter AgO treatm ent, removed Ag"*" with PO^, no Agi v is ib le in KMnO^ so lution 1.64 3.28 4.10 23. A fter AgO treatm en t, removed Ag* w ith NaCl, recovery no good here ---- 24. Left-over solution 0 .1 1

♦Solutions 11-15 were previously treated with AgO. 56 a 1.00 ntl. delivery volume. The ether sençiles were within 0.006 ml. of 1 .0 0 ml.

The 1.00 ml. graduated pipette was calibrated with mercury.

Thi 8 pipette was used to add the potassium permanganate to the solutions from vdiich the 1.0 ml. samples were withdrawn. The pipette delivered approximately 1 percent higher than the observed readings. This correc­ tion was applied to the calculated molarities in the tables. Nearly all of the above solutions were made up to 6 ,0 0 ml. from which 1 .00 ml. samples were withdrawn. The molarities were calculated accordingly.

The 10.00 ml. buret used in titrating the samples was also cali­ brated with mercury. The error in this buret was found to be negligible.

Conclusions from the Recovery Data in Tables 1, 2, and 5

It is apparent from the tables that the recovery of permanganate in phosphoric aoid solutions to at least 5.1 normal in the acid was reproducible. In very concentrated solutions, the recovery was low.

(The iodine method was used in a ll sam ples.)

In all but very dilute solutions of hydrochloric acid, the recovery of permanganate was questionable. Certainly if the solution was allowed to stand for a period of time, as indicated in sample 14, the recovery was low. This is true in solutions as low as 5 normal in the acid, smd most lilrely in more dilute solutions than 3 normal. Possibly chlorine is liberated by the permanganate ion in a manner similar to the libera­ tion of iodine but at a much lower rate. If so, the chlorine does not remain as such since upon the addition of iodide ion, the chlorine would liberate iodine and thus cause no apparent loss of permanganate. In any 57

«Tent» th« use of hydroohlorio aoid In recovery procedures should be avoided when pemanganate ion ie involved.

In perchloric aoid solutions, the recovery of permanganate to at least 5*IN In the acid was quite reproducible. In many of the samples, additional aoid (sulfuric acid) was added to the sas^le with no apparent loss In permanganate Ion. The exact amount of aoid in the solution is not critical, but a point is reached in the more con­ oentrated aoid solutions where the permanganate ion begins to decompose resulting in low recovery values. To avoid this loss, the concentration of aoid should be kept near 3 molar or lower.

It is also seen from Table 2 that the perchloric aoid solutions treated with silver oxide give results which agree quite well with those not treated with silver oxide. It is, therefore, concluded that the presence of the silver iodide precipitate in the solution does not appreciably effect the results. The addition of enough sodium hydroxide to a permanganate solution to make it basic, however, does cause the results to be low. Filtering through a glass sintered filter caused negligible decomposition.

The results from the sulfuric aoid solutions are similar to those obtained in the perchloric acid solutions. As seen from these experiinents, the recovery of permanganate is reproducible in solutions at least

4 molar in sulfuric aoid. The addition of sodium hydroxide to these solutions does not appear to oause low values as it does in the perchloric acid solutions. The silver oxide treatment apparently does not effect the results. As mentioned in the Experimental Procedures se

section, the end point in the presence of silver iodide is not as

sharp as In a pure solution* However, an attem pt to remove the

precipitate In any form (with PO^, Cl*- sto.) causes much more error than

that due to a less sharp snd-point.

The conclusions therefore are that,

1. Permanganate ion can be quantitatively recovered in phosphoric

aoid, sulfuric sold, and perchloric acid solutions and possibly very

dilute hydrochloric acid solutions.

2. The ooDoentrations of the first three acids can be at least

5 molar.

5* The silver oxide treatment does not appreciably effect the

recovery r e s u lts . I t should be remembered here th a t th e siIv e r oxide

ie simply added to the solution which is then heated and oooled and

then titrated. No oxidation occurs in the manganese since it is

already in the +7 state.

Recovery Date from Mn** Solutions and Mn**-MnO^ Solutions

Table 4 below gives the r* covery data for manganese in the +2

oxidation state using the silver oxide method described in the Experi­ mental Procedures sectio n . These so lu tio n s contained only manganous

sulfate which was subsequently oxidised to the +7 state with silver

oxide.

The volumes of the above solutions were approximately 5 ml. end

contained from Ifc to SO drops of concentrated sulfuric aoid. The

observed value of the molarity agrees with the calculated value within 59 experimental error. The latter value warn obtained from the weighed quantity of manganous sulfate dissolved in the original solution. It appears from the data that a quantitative recovery of the manganous ion fraction can be made to within 1 p ercent.

Table 4

Recovery of Manganous Ions

Sample Ml. 0.05 M Ml, (0.0099N) Observed Calculated Number MnSO^ Used *** 2®2^ Molarity x 10^ Mol. x 10^ 1. 0.085 2 .10 5.01 6 .0 1 2 . 0.085 2 .10 5.01 5.01 3. 0.083 2 ,12 5.08 5.01 4. 0.083 2,07 4.94 5.01

S. 0.083 2 .12 6.05 5.01 6 . 0.085 2.11 6.05 5.01 7, 0.083 2 .1 2 5.05 6 .0 1 8 .^ 0.085 2.20 6.26 5.01

♦Incomplete removal of AgO.

Recovery data fror both fractions in the Mn‘*"*^-MnO~ system are given in Table 5 on the following page.

Each sample in Table 5 contains 1 ml, of the original starting solution (6.00 ml. total). The total number of mi 111 liters of C.01Û0N sodium thiosulfate required for both fractions, after pipette correc­ tions, is 14.45 ml. This means that the total conoentration of each sample should be 4,82 x 10”® M. The main deviations from this value are due to

1. Reduced oxidation states of the persianganate ion which

were not recovered in the +7 state.

2. Error in the delivery of the volumetric pipettes used. 40

TabU b

++ EeooTvry Data on tha Syateai Ma

Sampi# la* fotto. WO^ Ml. Ma2®2Û5 Coao. Ma+^ Total Time (mln] fra o tlo a fra o tlo a Ma^*fraotlon fra c tio n Cono. Elapsed X 10* M /l X 10® m/ 1 X 10® Run V i 1. 1.81 5.62 0 .6 8 1.16 4.68 16 2 . 1.78 5.66 0.65 1.50 4.78 29 3. À. 1.74 5.48 0 .6 6 1.50 4.68 48 4. 1.76 5.52 0 .6 6 1.52 4.74 65 6. 1.66 5.52 0.75 1.46 4.68 75 6 . 1.59 5.18 0.80 1.60 4.68 91 0.26 ml. loft over from the above ruB 7. 1.74 3.48 0.61 1.22 4.70 15 8 * 1.72 5.44 0.64 1.28 4.72 29 9. B. 1.68 5.56 0 .6 6 1.32 4.68 49 10. 1.61 5.22 0.69 1.58 4.60 63 1 1 . 1.53 3.06 0.75 1.46 4.52 76 12. 1.26 2.62 0.83 1.66 4.18 91 0 .1 2 ml. left over from the above run 13. 1.80 3.60 0.55 1.06 4.66 16 14. 1.80 3.60 0.46 0.92 4.52 29 15. C. 1.77 5.54 0.61 1 .2 2 4.76 49 16. 1.72 5.44 0.63 1.26 4.70 62 17. 1.70 5.40 0.61 1 .2 2 4.62 77 18. 1.64 5.28 0.61 1 .2 2 4.50 89 0.18 ml. left over from the above run 19. 1.81 5.62 0.59 1.18 4.80 16 2 0 . 1.79 5.58 0.61 1 .2 2 4.80 51 21. D. 1.77 5.64 0.65 1.90 4.84 46 2 2 . 1.74 3.48 0 .6 6 1.32 4.80 60 25. 1.67 5.54 0.73 1.46 4.80 75 24. 1.56 5.16 0.83 1 .66 4.82 89 0 .1 2 ml. left over from the above run 26. 1.65 5.30 0.58 1.16 4 46 16 26. 1.65 3.50 0.62 1.24 4.54 51 27, E. 1.59 3.18 0.65 1.50 4.48 46 28. 1.48 2.96 0 .6 8 1.36 4.52 60 29. 1.56 2.72 0.75 1.50 4.22 76 30. 1.15 2.30 0.84 1.68 5.98 89 0.13 ml. left over from the above run 51. 1.75 5.46 0.60 1.20 4.66 16 92. 1.70 5.40 0.60 1 .2 0 4.80 51 55. F. 1.72 5.44 0.61 1 .2 2 4.66 45 34. 1.68 5.36 0 .6 8 1.56 4.72 60 55. 1.57 5.14 0.71 1.42 4.56 75 56. 1.52 5.04 0.80 1.60 4.64 89 41

S* Experimental error «hioh aocouats for approximately 5

percent of the total error obserred.

In run A, the values of the oonoent rat ion of Mn** (4th oolumn) are low. This was apparently due mostly to the pipette used («hioh delivered approximately 0.975 ml.). In plotting the oonoentration of the manganous ion frao tlo n verses tim e, the values «ere oorreoted to a

1 .0 ml. delivery. (This graph «111 be discussed after the other experiments are disoussed.)

The total volume of sodium thiosulfate used in both fraotions of run A is 14.59 ml, oonpared with the oaloulated value 14.45 ml. This is «ell within experimental error. Both fractions of run A «ere treated with silver oxide. Run B «as identical to run A except the permanganate fraotlon «as not treated with silver oxide. The values in this fraction are all low indicating the presenoe of lower valenoe states of aianganese. Howevr, samples 27 and 28 in Taole 2 on page 33 shov. that the addition of sodium hydroxide to a pure permanganate solution in perohlorio aoid results in low recovery values.

A study of the reaction MnO^ ste^ MnO^ has been mat' i i t wry basic solution (13.8M potassium hydroxide) and it is found tlc ' ri' first step occurs rapidly compared to the slower second step.

< equations and pseudo-first order rate constants are given in £2 chv roferenoe. The faot that permanganate does noticeably decompose even in dilute basio solutions, plus the faot that the isotopio exohange reaction observed in hydrofluoric aoid solutions is interpreted to indicate complete removal of the intermediate +3 valanoe state from

^^Jezowska-Triebiatowska, B, et al, Roosnlkl Chem. 29, 259-76 (1955). 42 the permeng&a&te fra o tlo n aa Indicated by the very low a o tir ity obserred in this fraotlon; (this will be disoussed later), would lead to the oonolusion that the intermediate ralenoe states (l.e. oxidation states of manganese less than the +7 state) in the permanganate fraction are +6 and +5 (mostly + 6 ) due to the sodium hydroxide present in the solution, and praotioally no +5 or +4 states. Thus, if the latter oxidation states are removed, then the isotopio exohange observed in the above system Ih** -MnO^ must be real slnoe rapid eleotron transfer 3- ** m — appears to ooour between species suoK as MnO^ and MnO^ -MnO^

(see page 6 of Introduction).

The OH ion Is the reduolng agent in the reaction UnO^ —^ MnO^.

Oxygen is erolred as one of the products*^ Slnoe the oondltions in run F were the same as those in run B, no silrer oxide treatment in the permanganate fraction, the same reasoning oan be used here for the low reoovery values. The experimental procedure for runs C, D, and E was the same ms for run A. The erratic values obtained in the manganous ion fraction in run C oannot be explained at this time except to say that the reoovery procedure for this fraotlon was not followed as outlined. The low values obtained in run E in the permanganate fraotlon were due to allowing the saïçles set for a day before titrating them.

A precipitate had formed in the samples whioh would not dissolve even when treated with silver oxide. The total volume of sodium thiosulfate used in run D is 14.55 ml. as oompared to the oaloulated value 14.45 ml.

Again the agreement is within experimental error. The same is true for run C except for the last two manganous ion samples.

^^C arrlngton and Symons, J. Chem. Soc. 1957, 659-65. 43

Lxperlmsnts on the study of the décomposition of permengenate solutions in the presence of rerioua concentrations of sodium t^droxide shows that in a SO minute period 8 percent of the permanganate had decomposed in a 0.115M NaOH so&ution, 10.2 percent In a 0.210M solution and 12.4 percent in a 0.262M solution* Since in the isotopic exchange reactions, the final permanganate fraction was approximately

O.ION in sodium hydroxide, the low values obtained for this fraction when not treated with silver oxide would be due to the reaction of permanganate with the base rather than incomplete removal of the lower oxidation states from the manganous ion fraction. Further evidence of this is the fact that the rate of decomposition of permanganate after removal from the manganous ion fraction corresponds much closer to the rate in the presence of sodium hydroxide than in the presence of the intermediate oxidation states as determined from spectrophoto- metrio data. These data are given in a later table.

The most conclusive proof, however, is from the isotopio exchange runs in hydrofluoric acid solution idiioh showed complete removal, within experimental error, of the +3 state idien sodium hydroxide was added. This is shown in the next section. To supplement this proof is the fact that the amount of smnganese preoipitated increases as the run progresses which indicates precipitation of the intermediate states.

This is shown on the graph in Figure 3. À further discussion concerning the graph will be given at this time.

Although the p o ints from the runs are somewhat soatt e red, the majority of them fall within experimental error. The solid line represents the average values through all of the points. .70 -

60

.50

40 e

> 1.30 (BC # o * 1.20 o ** I.IO o % ^ 1.00

0.90

0 10 20 30 40 50 60 70 80 90 100 Time (Minutes)

Figure 3. Concentration of Mongonese in precipitate formed by neutralizing solution with Sodium Hydroxide at various time intervals from the system Mn"*"*" - MnO^ — HCIO 4 • 46

The signlfioeaoe of the T&lue of the lateroept et time sero ahould be noted here. The oonoentration of meoganeee et this point is

1.16 X 10~^M. Assuming the Guyerd reeotion is taking pleoe in all of the solutions, then

2 MnO^ + 5 Ifa** ♦ 2 HgO == 5 MnOg + 4 H* (Guyerd R eactio n )

Thus for every 3 moles of manganous ion that react, 2 moles of perman­ ganate also reaot. The oonoentration of permanganate used in the above experiments was 4.10 X10"^M. and the manganous ion oonoentration was

6.9 X 10~^M. Assuming the reaotion goes to oosipletion, whioh is the case in excess permanganate solutions,^ then in the absence of inter­ mediate oxidation states (time sero) the oonoentration of manganese in the precipitate should be 6.9 X 10“^ plus 2/3 X 6.9 X 1 0 or 4.6 I 10”^M or a total of 1.15 X 1D“®M. This Is preoleely the intercept value,

(Figure 3).

The sharp rise in the curve towards the end of the experiments can be explained as due to the presence of a precipitate in the samples before the sodium hydroxide was added. This precipitate could usually be detected within 60 minutes after a run was started. DATA AND RESULTS PROM THE ISOTOFIC EICHANBE EXPERIMENTS

In this aootioa the date mud results mre given first for the experlmeats in solutloae of perohlorio mold. This mill include exp^rl- ments in ehioh sodium hydroxide was used to separate the two fractions and runs in sdiioh Dowex 1-X8 resin was used for separation. These data and results are followed by the experiments made In sulfuric acid and hydrofluoric acid respectively. Resin separations, if any, are include^ for these acids also.

Data for the perohlorio mold experiments are given in Table 6 as indicated therein. (Table 6 is on page 51.) A discussion of the equations used in treating these data and oorreotions applied are given below. This also includes a saisie calculation of the final results.

The equation used in the isotopio exchange experiments is given in the theory section as

/n (1-F) - -Kt where K is a constant isee theory section). The manganous ion fraction is used here to show how the fraction exoha age F was calculated. This latter quantity is defined as

F ■ S . 8 , Sqo *

T^ere S is the specific activity lAiloh is equal to the number of radioactive manganous ions divided by the total number of manganous ions in the so lu tio n . For a homogenous system

46 47

Ber* (à *) il the oonoaatratioa of r*dioaotlT* minganou* long Cgrmm atom p er liter) and (A) the total oonoentration of aanganoua ione both at tim e t . The equation for F then beoomea

T . ( ! ' ) - U ' o J - U ' , )

SubaorLpta aero and infinity refer to time.

If the aame oounting oryatal ii uaed throughout a given experiment and the aame geometry and oounting effioienoy la maintained for eaoh manganoua ion eai^le, then the oonoentration of the radioactive manga ne ae i* direotly proportional to the aotivity of the eample. The propor­ tionality oonetant would oanoel out, leaving

(Àot)^ - (Aot)g ^ " tAfft)oo - Uct)g

The quantity (Aot)^ refera to the activity at time t. This equation is exact if the separation procedure produces no exchange. In the prooess of separating the two manganese fractions, there is undoubtedly some

"separation induced" exchange, but the extent of this exchange suiy be difficult to determine. This will be disouased at the end of this s e c tio n .

In oDtaining the value of (Aot)^, a small oorreotion was first applied to the observed oounta of eaoh fraotion. This oorreotion was believed necessary because of the deoon^osition of the permanganate in solution as the experiment progresses. A sample oaloulation is as follow s* 48

Assme the obaarred activity of tha penaaagftnate fraction after

IS fldnutea ia 500 cpm, and 1000 opn in the manganoua Ion fraction alao

at 15 minutea* From Figure 4, (page 6S), vdiioh maa obtained from

Figure 5, the permanganate concentration at this time la 5*61 I lO'^M.

Subtracting thia from the intercept value 3.64 I 10"®M (t “ O) glvea

the amount of permanganate deoompoaed» 0.05 X lO'^M. The percent

recovery value la 5.61/3.64 or 99.2 percent. Converting to a 100 percent

recovery basia, the cou at a/ml nut e for the permanganate fraotion ia

y ■ 600/0.992 - 504 cpm

Applying the correction now to the manganoua ion fraotion gives

X - 1000 - (604 - 500) ■ 996 cpm

Thia procedure waa uaed for calculating the activity of both fractions

of all a amples in idii oh sodium hydroxide waa used aa a means of sepa­

ration. These are the corrected actIvitiea listed in the tables. The

corrections are more important in later samples, e.g. after 46 minutes

instead of 15 minutes.

The value of (Act)^ was obtained from the sum of the activities

of both fractions from any one sample. This value ahould be the aame

for all samples in a given run in which no precipitate appears before

the addition of sodium hydroxide. Low values of this sum indicates

the appearance of a precipitate. At the concentrations used in these

isotopio exchange runs, the precipitate usually appeared within 60

B iinutes. 49

The quantity (Aot)^^ oan be oaloulated from the value of (Aot)^ and the initial oonoentration of the reaot ante in the solution which in this oate would be potassium permanganate and manganous sulfate.

A general equation for caloulating this quantity is

This is for the manganous ion fraotion and (A) and (B) are the total msmganoue ion and permanganate ion oonoentrations respeotirely. The same quantity for the permanganate ion fraotion is obtained simply by replacing the (A) in the numerator of the above equation with (B).

Of oourse the above aotlritles are oaloulated on the basis that no décomposition of permanganate has occurred. Speotrophotometrio data shows that a small but detectable amount of the permanganate does de- oompose over a period of 60 minutes, but the eurrature this would produce in the ^n (l-F) verses time plots oan hardly be detected.

In addition to those runs made in an atmosphere of air, other experiments of iaentioal ohemioal composition as those siade in air were made in other atmospheres. For example, experiments v.cre made in an atmosphere of nitrogen and in oxygen, and finally an attempt was made to follow the exchange in an atmosphere of hydrogen. The effect of these different gases on the rate of the isotopio exchange is seen in the half-life values obtained from graphs from Figure 6 through Figure 15. 5 0

In the experiments in e nitrogen atmosphere, the gas was bubbled through the solution for approximately lb minutes prior to the addition of the potassium persianganate. After the latter was added, the gas was bubbled through the solution for a few more minutes and then the gas was passed over the solution during the entire run in order to keep an atmosphere of nitrogen present at all times* The same pro­ cedure was used for the other gases also. The data and the results are given in Table 6 as indicated in the parenthesis after the letters

(a , B, C, etc.) representing the experiments, and the graphs are given in Figures 11 and 12.

Discussion of the Results Obtained from the Mn** -UnO^ Isotopic

Exchange Reaction

The results obtained from this isotopic exchange reaotion are given in the tables and grsipha on pages 51 through 56. It is seen from the graphs that the plot of (1 - F) verses time on semi-logarithmic graph paper does not give a line that passes through the ordinate at 1 when the time is equal to zero for some of the experiments. This is typical of exchange reactions in which there is a separation-induced exchange. (A separation-induced exchange is an exchange, in excess of the normal exchange, caused by the separation procedure, eg precipitation of manganese dioxide in the Mn*^-MnO^ system.) It was stated on page 47 that the extent of this exchange may be difficult to determine. Actually, the difficulty was thought to lie in determining the value of the activity at time zero, Wiich is called the apparent zero-time exchange. The reason for thinking this determination would be difficult was that the value of the quantity (Act)^, used in the T able 6

Isotopio Exchange Data

P ercen t Corrected Activity Corrected Activity Sample Recovery t 1/2 Permanganate Ion Manganous Ion Time Number Run la o id ) Permanganate (min) Fraotion (cpm) Fraotion (opn) (1 - F) (Min.)

1. A (BOIO4) 98.2 749 3632 0.805 28 2. 96.6 966 5425 0.747 47 (Act)^ ■ 4370 cpm 165 5. 94.1 1171 3195 0.686 64 4 . (A ct) 00 - 3740 cpm 89.3 1178 3167 0.678 85

1. B (HC104) 98.2 793 5691 0.791 28 2. 96.2 1005 5370 0 .732 50 (A ot)o 4376 cpm 3. 95.5 181 1162 3277 0.707 67 4 . (A o t)^ 3745 cpm 87.9 1267 5088 0.656 86

1. C (HCIO^) 99.0 1040 5629 0 .825 16 97.8 1456 0.760 2. (A ot)^ 6620 cpm 168 6261 36 3. 96.6 1674 4904 0.697 56 (ACt)gp 5668 cpm 4 . 9 2.1 1869 4TO7 0.6 6 2 74

1. D (HCIO^) 98.2 1558 5476 0.780 28 2. 9 6 ,7 1652 6154 0.721 45 (ACt)g 6740 optm 160 3 . l A r t ) ^ 5760 opn 94.0 1782 4941 0.688 64 4. 90.1 1881 4746 0 .664 81

1. c t 0.428 18 This is a reoaloulation of run C given above. The same values of the 0.389 56 activities are used in the oaloulations here as recorded above, except 0.561 66 that these activities are divided by a oonoentration term. The values 0.550 84 tabulated here are ( 1 - T) ; F is defined on page 59» oi Tablo 6 Contd.

P eroent Corrected Aotivlty Corrected Aotivlty Sample Recovery t 1/2 Permanganate Ion Manganous Ion Time Number Run ( a d d ) Permanganate (min) Fraction (opn) F ra o tio n (opm) (1-F) ( la n .)

1. 0 (BCIO^) 179 1299 0.880 16 2 . (Aot)^ ■ 1450 opn 156 341 1160 0.780 39 5. (A ct)o ^ " 1250 cpm 329 1079 0.704 66 4 . 438 ♦ 50* TOI + 260* 0.600* 108

1. H (HNO3) 996 30,997 0.963 7.8 2. (Aot)o ■ 32,000 opn 1843 30,218 0.9 3 5 14.5 5. 168 2960 28,905 0.887 29.5 (Aot)op" 27,560 opn 4 . 3179 28,660 0.878 37.8 3. 3661 28,686

1. I (HCIO4 ) 964 29,964 0.963** 0.974 7.5 1617 28,975 2. (Aot)^ - 30,660 opn 169** 0.942** 0 .9 3 6 13.5 5. 2627 27,893 0.900** 0 .895 21.5 (Aot)jjp- 26,200 opn and 3020 27,869 0.885** 0.893 29.5 4 . 182 6. 5677 27,134 0.880** 0 .8 6 8 37.5 6 . 4794 26,842 0.817** 0.816 4 5 .8

1 . K (BCIO4 ) 99.1 569 3740 0.856 23 5666 41 2 . (Aot)^ ■ 4240 opm 98.5 728 0.806 97.4 213 899 3288 0.725 61 5. (Aot)oQ" 3468 opn 4 . 95.3 954 3256 0.710 84

* Correction applied to low values ia this sample. These are 1 -F values from the permanganate fraction. All others are from the manganous ion fraotinn.

on T able 6 Contd.

P ero en t Corrected Aotlelty Corrected Aotielty Sample Re 0070 ry t l /2 Permanganate Ion Manganoua Ion Time Humber Run (a c id ) Permanganate (min) Fraotion (o^m) Fraction (cpm) (1-F) (M in.)

1 L (HCIO^) 99,0 588 5877 0 .8 4 6 26 9 8 .1 844 5661 0.786 47 2 (Aot)^ ■ 4440 opm 5 (Act ) oo " 5652 opm 96.4 218 957 5462 0 .7 5 1 71 4 94.2 1015 5401 0.713 94 5 9 1 .4 1067 5525 0.6 9 5 114

1 M (HGIO^) 99.7 2892 29,577 0.904 11 99.1 5891 28,049 0.857 17 2 (Aot)(, " 52,000 opm 5 98 .4 119 4675 27,514 0.851 25 (Ant)QQ- 27,100 opm 4 97,7 5596 26,156 0.790 55 5 95,5 6866 25,319 0.768 41

1 0 (HCIO4+N2) 99.0 840 4760 0.825 18 2 (Aot)^ ■ 6600 opm 97.9 1021 4657 0.799 52 4529 0.766 47 3 (Act)go" 4790 opm 96.5 182 1166 4 94.7 1004 4186 0.706 61 5 9 1 .7 1652 5834 0 .6 5 1 75

1 P (HCIO4 +H2) 98.9 949 5527 0.821 19 2 (Aot)^ ■ 6290 opm 97.9 1256 5086 0.776 52 96.5 164 1462 4856 0.750 47 5 (Aot)o(j* 5580 opm 4 9 4 .5 1600 4768 0.715 62

01 «• Table 6 Contd.

P eroent Corrected Aotleity Corrected Activity Sample Reoorery t I /2 Permanganate Ion Manganoua Ion Time Number Bun ( a d d ) Permanganate (min) Fraction (opn) ______Fraction (opm) (l 11 (MLn,) 1. Q (BDIO^+Hg) 90.0 1065 6148 0.818 14 2 . (A ct) ■ 6150 cpm 87.4 1439 4755 0.746 50 5. 85.2 155 1609 4555 0.706 45 (Act)gp" 5500 opm 4 . 72.4 1661 4166 0.659 59 5. 65.2 2214 3194 0.403** 72

1. R (EClO^+Og) 98.5 945 6425 0.8 2 5 16 2 . 96.1 1512 5114 0.768 29 5. 94.5 1505 4807 0.712 40 (Aot)^ ■ 6580 opm 4. 92.0 104 1691 4752 0.698 60 6 . (Aot)ojj- 5460 opm 88.4 1697 4660 0.6 8 5 60 6 . 80.0 1876 4545 0.664 71

1. S (H2SO4) 99.5 4645 28,587 0 .8 4 7 9 2 9 9.1 5449 27,708 0.814 16 . (Act)(j ■ 52,900 opm 5. 96.6 157 6916 26,780 0.782 25 (Aot)gg " 28,190 opm 4 . 90.7 6424 26,586 0.768 51 5. 81.1 7700 25,205 0.726 40

1. S (HF) 445 31,545 0.984** 0.985 10 2. 815 31,057 0.970** 0.969 26 ( lo t) , 51,900 opm 5. n e g lig ib le 568 51,203 0.978** 0.975 55 (A ct) 27,280 opm 4. 00 684 51,595 0.979** 0.989 41 6. 821 50,950 0.970** 0.965 61 3 .6 5

360

3.55

3.50

X 3.45 i 3.40

ç 335

3.30

325 0 10 20 30 40 50 60 70 80 Time (Minutes)

Figure 4. Decomposition of Potassium Permanganate in presence of Manganous Sulfate. (KMnQ,)= 4.1 x 10"^IVI, (MnS0^)=6.9% IO-^M.,and (HCIO^)»3.1 M. 5 6

(KMnQ,)-4.IO x IC^M

( 0 ) (MnS0^)-6.9 X I0"^M (HCI04 ) - 3 .l M

(KMrO^)-4.lOx 10"^ M (b) (MnS0*)-6.9 X lO'^M (HCIO* ) - 3 I M Figure 5. (a) ti = 163 Minutes ; (b) t±= 181 Minutes ______1______I_____ : ______L_

(KMn04)»4.l0x lO'^M (MnSQ4.)- 6.9 x lO’^M (HCIO4 )-3 .t M

(KMo0 4 )- 4 .I0 xIO"'M (M nS04)-6.9 X 10"^M (a')ti = 1 4 2 Minutes; (HC1O4 ) - 3 .i m Figure 6 . (a)ti = l58 Minutes; (b)ti = 150 Minutes ______L______I I______

(KMn04 )- 4.30 x IO’-'M (MnSO* )-6 .9 X I0"*M (HCIO4 ) - 3 I M

(KMnO^)-4.IOx IO"M (MnSOe )• 6.9 x 10“^ DOWEX I-X8 Resin used (HNO, )-3 .l M

Figure 7 (a)ti = 156 Minutes; (b) k= l 6 8 Minutes ______L______I______I------2 6 2 5 52.50 78.75 I05j0 Time (Minutes) 5 7

(KMn04)»4.l0x IQ-'m (MnSQ<)-6.9 x 10“ '* M (HCI04)-3.I M

(KMnO#)»4.IOxIO"^M Ü. (MnS0^)"6.9 X 10“ ^ I Figure 8. (o)MnCV fraction,t^=l69Mlnutes (b^ Mn*'*'froction^t^=162 Minutes _

TO

(KMnO*)-3.tOxlO"^M {MnSCU)-69x IO~*M (HCIO* )-3 .l M

(b ) (KMnO*)- 3 10 X IO"*M (M nS0*)"69 X 10""M

Figure 9. (a) tj=2l3 Minutes; (b)t^=218 Minutes

J______I______L iOÇ

(KMnO^)«4.IOxlO"-*M (MnSOr)-7.5 x 10"^M (HCIO4 ) - 3 .l M

(b) Ü. I Figure 10. (o) MnO# fraction; t^ = ll9 Minutes (b) tj(=163 Minutes; plot of first 2 points of ^ ______runs A,B,C,|Ond D.

n.5 3f S9..5 to

Time (Minutes) 9 8

( 0 ) (KMnQ»)-4 .IOxKT*M (MnSGU)-6.9xlCr^M (HCK)4)-3.IM

(KMnO#) «4 .10x 10- - Experiment# mode In Nitrogen AtmoephereCMnSQ*)« 6 .9 x 10*^M (HCKV )-3 .l M Figure II. (a) t^=l82 Minutes; (b) y= 164 Minutes ------1______I______I______

(KMnO*)" 9 9 6 x KT'^M (Mn9 0 e ) - 6 9 x IO*^M (HCK)4)-9I M

(g ) Nitrogen Atmosphère 0

(KMnQ4)«4.IOx 10' (MnS 0 #)-6 .9 X I0 "^M (H^SQ* ) " 3 .l M Experiment mode in Sulfuric Acid

(b) (KMn0 4 )- 4 .l0 X lO-^M Z (MnS 0 4 )- 6 .9 X 10" ''M (HF) " 3.1 M Experiment mode in Hydrofluoric Add Figure 13. (a) ti=l37 Minutes; (b) f^- «» ______I______I------1------17.5 3 5 5 2 .5 70 Time (Minutes) 60 oaleulatlon of F, would o»u## a eonalderablo error in the value of

(l - P) at time lero. This error would result from using the equation for F whioh is valid only if no separation induced exchange ooours*

Thus, the value of (Aot)^ used in the equation for F should be that of the pure manganous ion fraotion prior to separation* This, of course, is not the case. The value of CAot)^ used, most likely. Includes activity from both fractions. Therefore, the problem is to find out if the value of the half-time, obtained when these values of (Act)^ are used, is the oorreot value (within experimental error). A check was made using a modified equation for F idiich takes into account the separation induced exchange occurring In this system. The modified equation is

« ^ ~ Sq (Aot)^(MnOg) - (Aot)/(Mn**) “ s,,- 6. rrrr-TTtrr?------o (Aet)^y(Mn++). (Act)y(Mn++)

In the original equation for F, the quantity S ia the specific activity of the manganous ion fraction prior to separation. The equation for

S i s

g . 0 ^

(Mb**)

The (Un'*^) oonoentratlons cancel out leaving the equation for F as given on page 47. {

Caloulating the values of F and (l - F) and plotting them on the graph in Figure 6 (aO gives a straight line nearly parallel to the line obtained from the original (1 - F) verses time plot. The intercept is quite different but the resulting half-time value is within experimental error. The apparent sero-time exchange oan be calculated from the (l - F) or (1 - T) values at time sero. Thus, eo

800 - So ° ®oo - So

The epperent mero-tlmie exehange aetlrity is obtained from S in the equation for and from g in the equation for The calculation valuee of this activity from both equations agree within experimental e r r o r . It is seen from the plots of (l - F) and (I - T) verses tisks in

F ig u re 6 (a) that the main difference in the two quantities is the intercept value. However, since the two lines are nearly parallel, it is apparent that the two quantities differ only by a constant. Thus, the values of the half-time obtained from the (1-F) plots are valid within experimental error.

In view of the errors due to (1) the separation procedure used,

( 2) the presence of a precipitate in some of the samples before the addition of sodium hydroxide, (S) the teiiq>erature ohanges, (4) plotting the graphs, and (5) the calculated value of (det)^, the value of the half-time cannot be aiy more accurate than t-l / 2 ■ 160+50 minutes for these perohlorio acid experiments.

Data from the Dowex 1-X8 resin experiments also show that the half- time value is close to 160 minutes. In these experiments there was no attempt made to correct the activity of either fraotion. The values of

F were simply calculated from the observed activities of each fraction.

Figures 7 and 8 show the results of these plots* In run G, Figure 7 (a), the resin oolusms were long and the solutions were passed slowly through the colusms. In runs H and I, Figures 7 (b) and 6 , the columns were very short (less than three inches) and the solutions were passed 61

quioklj through tb# eolunaa* Th# liquid coming from tho oolwmmm mm#

porfootly olomr indioating r#moT#l of th# potm##iv#a pormmngmnat#.

It i# apparent from th# graph# that if the aolution i# pa###d

#lowly through th# #olumn, a aeparatlon-induced exohang# omoura, that

la, the line doea not paaa through 1 at time aero. Eommver, if the

aolution ia paaaed quickly through th# column, there appears to be no

aeparation-induced exchange* the line paaaee through 1 at time zero.

Thia can be explained if it ia aaaumed that the reain acta aa a oatalyat for the Guyard reaction or deooatpoaea the potaaalum permanganate. It waa found that the later prooeaa doea occur, but the reain may alao act aa a oatalyat * In either event, In order to get no aeparatlon- induced exchange, the prooeaa mu at be alow compared to the tixae it take# for the permanganate and manganoua ion fraction# to aeparat# when the aolution la paaaed quickly through the column. The alower the aolution ia paaaed through the column the more time the reain ha a to act and the greater the aeparatlon-induced exchange.

The Taluea of (1 - F) were calculated for both fraction# of experiment I. The reaulta from the permanganate fraction are plotted in Figure 6 (a)t thoae from the manganoua ion fraoticn are plotted in

Figure 8 (b). The half-time valuea obtained from the reain run# agree more cloaely than thoae obtained from the run# in adiich a odium hydroxide waa uaed to aeparate the two fractlona. From the reain rune, the half- time la calculated to be 164+16 mlnutea. The acid uaed in run H waa nitric acid. The calculated half-time agree# closely with the average v a lu e . 62

Tha ooDoaotratlon of potamgium parm&ngan&t# and manganou# tulfat»

aa# varied individually in #ome of the run# made i n perohlorio aoid.

For ezampl#, in the graph# #ho#m in Figure 9, it i# #een that the half-

time haa been extended from an average of 160 minute# to an average of

216 minute#. The rea#on for this is that the oonoentration of potassium

permanganate has been reduced from 4.10 X 1 0 M to 3.10 X 10“^ M.

Similarly, In Figure 10 (a) the oonoentration of manganous sulfate ha#

been inoreased from 6.9 X 10”^ U to 7.5 X 10"^ M. The value of the

half-tim e ha# been reduoed from 160 minute# to about 120 m in u te# .

Hovevar, this decrease in half^im e appears to be too great oompared

to the small inorease in manganous sulfate oonoentration. Nevertheless,

the value of the half-time is changing in the proper direotion. À#

mentioned earlier, experimental error# are rather large when using

sodium hydroxide as a means of separating the two fractions.

It appears in most of the graphs that the point# begin to deviate

from a straight line after a period of about 60 minutes. The last point

or two of eaoh experiment was not plotted simply beoause the deviation beoame greater and greater after the 60 minute period. The reason for this is that a preolpitate is forming in the solution thereby causing

the manganous ion fraction to inorease in activity rather than decrease

as it should. Therefore, the value of F is lower and (1-F) is higher.

The points always deviate in this direotion in runs where sodium

hydroxide was used for separation.

Since runs A, B, C, and D have identical chemical oompositions, and

since some of the sauaples were withdrawn at different time intervals, a

plot of ( 1 - F) verses time was made of the first two points of these runs 68 to detervin* how th# h#lf-tlm# obtained from th is plot oo^par## with that

of the individual ruiu. In taking th# flr#t two pointa, th# time would b# lea# than 60 mlnutea, therefor#, no preolpitate ahould be preaent In these sample a. The value of the half-time la 163 minutes whioh agrees with the average value of 160 mi mites. It is realised, however, that

In plotting all the graphs, the first three points In eaoh graph were

given more weight than the others. Therefore, it ia not too surprising that the 163 minute half-time was obtained in the plot mentioned above.

Is a final test on the rate of sxohange of manganous ion with

permanganate Ion in perohlorio aoid solution, several experiments were made In an atmosphere of pure nitrogen. An experim ent was a l s o made

in an atmosphere of pure oxygen. Hydrogen gas waa alao used, but the data from this run seemed to Indloate that the permanganate in solution

rapidly deoompoaed. The evidenoe for this is in the value of the total aotivlty of eaoh sample withdrawn. Addition of the activity in the manganous ion fraotion to the activity in the permanganate ion fraction

should give approximately the same value for eaoh sample withdrawn.

Thia is true for at least the first three or four samples in each of the other experiments, but the experiment in an atmosphere of hydrogen

showed a oontinuous decrease in the total aotivlty for each sample from the very start of the run. The decrease was too great to make the data reliable.

Within experimental error, the résulta..from the experiments in an atmosphere of nitrogen give the same value for the half-time as the results from the experiments laade in an atmosphere of air. Figure 11 (a) 64 and (b) and Figure 12 (a) are the graphs for the expérimenta in a nitrogen atmoaphere. The reaaon the half-time value from Figure 12

(a) la leaa than the average la beoauae the permanganate oonoentration waa inoreaaed from 4.10 X 10"® M to 5.86 X 10"® M.

In view of the above reaulta, nitrogen appeara to have little or no effeot on the rate of the laotopio exchange» If the gaa la bubbled through the aolution prior to the addition of potaaalum permanganate, and an atmoaphere of nitrogen la kept over the aolution during the run, then the appearanoe of a precipitate of manganeae dioxide aeems to be delayed aevera 1 mlnutea oompared to the time it takea for the preolpi­ tate to appear in a aolution open to the air. However, the reverse la true when an atmoaphere of oxygen ia uaed* The preolpitate aeema to appear earlier. The value of the half-time la alao reduoed from approximately 160 mlnutea to about 104 mlnutea aa indicated from the graph In Figure 12 (b). Thua, the addition of oxygen gaa to the aolution doea effeot the rate of the laotopio exohange In the Mn**-MnO^ ayatem. i^parently hydrogen gaa alao has a very noticeable effeot on this system.

A possible explanation of the oxygen effeot on the rate of exohange in the Mn - IfnO^ ayatem la that the reduolng power of the oiqrgcn oxidises the +2 and +5 oxidation states of manganese to the next higher oxidation states. It would seem from the reaulta of the e^qperimenta, t h a t th e *2 state la affooted laore by the oxygen gaa than the +3 state.

The reaaon for believing this la that a preolpitate of manganese dioxide did not appear in solution for at least 50 minutes after mixing. (Thia waa not the case when hydrogen gas waa used. See next paragraph.)

Possibly the oxidation of the +3 state la slowed down by the weak ooa^lex 66 format!oa la this state* That oxygen does oxidise the *2 state is erident from the slow oonrerslon of UnO to NbgOg and MnOg in the presenoe of air*^* Merertheless, with the appearanoe of higher oonceatrations of the intermediate oxidation states of smnganese in the solution, a faster rate would be observed since the intermediates are carrying the exchange process*

It was stated, on page 63 that from the initial time of ndzing, the permanganate in solution decomposed rapidly in the presenoe of hydrogen gas* If it is assumed that the permanganate reacts quickly with the hydrogen gas to form lower oxidation states of manganese, then, with the rapid appearance of these lower oxidation states, the in itial concentration of Mn^* would be high in this solution compared to a solution in an atmosphere of air or nitrogen* This would result in a more rapid appearanoe of a precipitate of manganese dioxide*

This was actually the case as was indicated by the decreasing total activity of each sample withdrawn* (See page 65*) Thus, it would appear that the rapid decos^osition of permanganate and the rapid appearance of manganese dioxide can be explained in terms of the reducing ability of the hydrogen gas.

No half-tim e exchange was determined from this experiment beoause of the rapid appearanoe of manganese dioxide in the solution.

The final experiments made in the Nn -UnO^ system were made In the presence of sulfuric acid and hydrofluoric aoid individually* The graphs for these experiments are given in Figure IS (a) and (b) respectively*

^mAlpine, R.K*, and Soule, B. A*, Qualitative Chemical Analysis, Page 546, New York, D. Van N ostrand Company Ino* 66

The h&lf-tlm# valu# obtained from the lulfurlo acid experiment 1#

157 minute*! * lightly lorwer than the average value of 160 minute*. It

i* possible that this lower value i* due to the fact that potassium

permanganate, In the presenoe of manganous ions, décomposés faster in

an aoid solution than in a perohlorio aoid solution. This is shown in

the speotrophotometrio studies given in a later section. This inorease

in rate of deoompositioa of potassium permanganate may be correlated

with the very slightly greater eosqplexlng ability of the sulfate ion

as oompared to the perchlorate ion. In general, the greater the oom-

plexing ability of th* aoid anion, the greater the rate of décomposition

of permanganate in the presenoe of manganous ion. The overall result

is an inorease in oonoentration of the intermediate oxidation states

of manganese. Unless a fairly strong complex is formed with one or more

of these intermediate oxidation states, this increase In oonoentration could aooount for an inorease in the rate of the isotopio exohange reaction. In the case of the sulfate ion, the complexes formed with the intermediate oxidation states are weak (oompared with fluoride ion, for example). (See the speotrophotometrio studies in later section.)

Therefore, the exohange prooess may continue. In the presenoe of these weak complexes, nearly as fast as if they were not present at all. If this is true, the overall rate of exchange wo-Id inorease as it appeared to do in the presenoe of the sulfate ion.

The fact that there sms no exohange observed, srithin experlsMutal error, in the ssuaple solutions containing hydrofluorio aoid oan be explained in terms of the oomplexing conditions mentioned above. The 67 fluorld* Ion form# strong oomploxo# with on# or moro of tho Intormodlmto oxidation state# of manganese a# indioated hy the very rapid deoompoaition of permanganate in the presenoe of manganoua ion and fluoride Ion. In the resin separations, mhioh are disouased in the latter part of the next seotion, it is shown that the fluoride ion is ooi^rlexlng with the +3 oxidation state of manganese in this system, j^parently this oomplex is so strong that any isotopio exohange beyond the +3 oxidation state of manganese is either very alow or does not ooour at all. Thia means that the fluoride oomplex ef the +5 oxidation state is not labile with any higher oxidation state of manganese. It is shown in the next seotion, however, that the complex is labile with the +2 oxidation state of manganese.

This concludes the study of the Mn** -MnO^ system in ehioh radio­ active manganese is used as a tracer. The next seotion deals with the oomplexing ability of aoid anione with the manganous and manganlo ions. The last seotion covers the rate of décomposition of perman­ ganate in the Mn** -ïlnO^ system.

It should be mentioned here, before going on to the next seotion, that the E value, as defined on page 5, was oaloulated for the perchloric aoid runs and found to be approximately 6.4 hours. The value of £ and the value of the half-time are somewhat lower than the values obtained by Adamson in his experiments. No reason for this oan be given at t h i s tim e . DEIERMINATION OF THE CONPLEXING ABILITY OF HANOANOUS IONS NITH

NOg, CIO4 , moj, HgPO^, F IONS AND TEN DI-SODIUM SALT OP

mTLENEDIAMlNNTETRAACBTIC ACID

The oh#mi 0*1 properties of lone in the vnrloue oxidetion etetee lends to the prediction that the mte-detemdning step in the Mn^^-lAiO~ isotopio exohangs may be between the hydrolised forms of the +3 and +4 oxidation states of manganese. Eoeever, there is the possibility that th e +2 oxidation state of manganese is forming a oonçlex with the above aoid anions, especially the HgPO^ and F ions* If the manganous ions are forming complexes with the aoid anions, then the rate-detem ining step could be between the +2 oxidation state and some higher oxidation state of manganese instead of from the +5 to the +4 oxidation states as indioated above. But in view of the fact that the +2 and +3 oxidation 2 & states of manganese in systems such as , 1*3 -lfei(aoetyactonate) g or

oxalate)g , exchange 00^ lately within time of separation, it would seem even less likely that the +2 oxidation state of manganese is involved in the rate-determining step. Nevertheless, it seemed desirable to study the complexing ability of the manganous ions, beoause the forma­ tion of a oomplex with these ions might still involve the possibility o f th e +2 oxidation state of manganese entering into the rate- determining step. The fact that complete exchange was found in the systesu mentioned above does not mean that the +2 and +3 oxidation states of manganese w ill always exohange oompletely or even at all.

^^Drehmann, D ., Z physik. Chem. 63B, 227 (1943).

68 60

la f&ot, the exohaage between theee two oxidation atates in a lolution

of fluoride ions wa# found to be inoonqplete even after one hour and ten

■dnxitei after mixing* Thic w ill be ditoueaed later along with other

reauIta obtained from reain aeparationa.

In the atudy of the ooaqplexing ability of the manganoua Iona it waia found that theae iona formed only weak oo^lexea with the following

iona, CIO^, NOg, 80^, HgPO^ and F . The iona are arranged in order of

inoreaaing ooaçlexing atability. The oomplex formed by ETTA waa found

to be muoh greater than any of the aniona liated above. A disouaaion

of the raaulta la given below.

Diaouaaion of the ReauIta Obtained from the Study of the

Coa^lexing Ability of the Manganoua Ion

The aaturated aolutlona of manganoua oxalate containing the

oomplaxlng ion remained in the tei^erature bath from 24 to 48 hour a

unleaa atated otherwiae. The prooedure uaed for atirring, reawving

aai^plaa, adding aalta etc. waa identical for all runa. Therefore,

unleaa the approaoh to equilibrium waa very alow, the reaulta obtained

give a meaaure of the relative atrength of the complexea formed with the varioua aniona. The EDTA oomplex formed with the manganoua ion waa 2g uaed aa a reference point. Thia ia a atrong oomplex. The log of the

Keq f o r t h i s oomplex ia 13.58.

The first experiments, in the atudy of theae complexea, were made in very dilute aolutlona of the varioua aalta used. The ionio strength

26 Sohwaraenbaok, 0., and Freitag E., Belv. Chem. Acta., 34, 1505, (1951). 70

of the solutions mss kept within epproxlmetely 0.011 units. In this

range, the Debye-Hiickel equation for very dilute solutions should be

applicable. The equation is applied es followst

The solubility product of aanganous oxalate is represented by

Ks - S^f^ - Sf

and S refer to the solubility of smnganous oxalate in pure water and

in the smrious salt solutions respeotlTsly; fg and f are the oorrespond-

ing mean aotisity ooeffieients. Taking logarithms gives

log S ■ log S^f^ - log f

Applying the Debye-Hüokel equation (see Theory Seotion) glves^^

log S ■ log SpÇ» + As+t_ J7T*

o At 25 C. and for water as a solvent, the value of A is about 0.50. Thus,

a plot of log S verses should give a straight line with a slope equal to 0.5t+s_ which for manganous oxalate is squal to 2.0 if no complex is

formed and the increase in solubility is due solely to 1uterionio attrac­ tion effects on whioh the premise of the Debye Limiting Law is devised.

Actually a plot of the data gave a variation in the value of the

slope from 1.40 to 1.76 (see Figure 14). The deviation from the calcu­

lated value of 2 .0 oould be due to at least two factorsj first, an equilibrium was not attained in the salt solutions at the time the sasq»les were withdrawn; second, the ionic strength of the solution was

outside the range in which the Debye-HUckel equation is valid. The

^^Prutton and Karon, "Fundamental Principles of Physical Chemistry," New York, The MaoMillan Cox^ai^, 1951, 0 NoNQa Sen# 035 A MoCiQ» Sen# V N«t9Q« San#

054

035

s 032

031

030 94 9ft 98 100 Ktt 104 lOft m Nt #4

10*

F i^ n 14. Sotubilrty of MnCjO^ in the preia nce of various soHt. Tt

latter ooodltion la beliaved to ooutribute mora to the deviation in the

■ lopea than the former. The reaeon for believing thia, ia that Unea drawn through the pointa in i^oh the ionic atrength of the aolution waa leas than 0 .0 1 0 2 gave a value of the alope oloaer to 2 .0 th a n d id the lines in irtiloh all of the points were taken into oonaideration. A gradual leveling off of the pointa waa obaerved as the ionio strength inereaaed. Both of the oonditiona mentioned above would oauae the points to fall below the straight line aa the ionic atrength inoreaaed. Those points that deviate considerably from a straight line at hl^er Ionio strengths were moat likely effected by both oonditiona.

Undoubtedly, some of the error observed in the low slope values waa due to inaoourate determination of the manganese in theae very dilute solutions, but in view of the faot that all values were below the expected value of 2 .0 would seem to Indioate that the factors mentioned in the paragraph above are quite important and most likely the predominant ones.

Regardless of which condition oausea the greater deviation, the important conclusion that can be reached from this study la that, over a 24 hour period, the solubility of manganoua oxalate in theae very low ocnoentrationa of theae possible complexing iona does not increase above that expected from the ionic atrength principle (within experimental error). Therefore, from the results of these experiments, it would appear that the manganous ion waa forming, at moat, only very weak coi^lexea if any in the salt solutions. However, the concentration of 71

Mingan«se wm# «lightly higher in the aolution oontaining fluoride iona

and in the aolution oontaining Iona, but the inoreaae in aolubility

ma not ooeparable to that obtained in the preaenoe of a atrong ooa^lex

auch aa manganoua iona with EDTÀ, Thia ia ahown in the following

expérimenta of high oonoentratione of complexing aniona.

The prooedure followed in theae expérimenta waa identical to that

uaed in the previoua expérimenta. The only difference ia in the aalt

oonoentrationa in the aolutlona* In the flrat expérimenta, the aalt

oonoentrati on waa kept very low ranging from about 0.001 to 0.004 mole a/

liter. In the following expérimenta, the oonoentrati on of aalt ma muoh

higher, ranging from 0.005 to 0.500 m olea/liter; In thia concentrâtion

range, the ionio atrength ia muoh too high for the Debye-Buokel equation

to be applicable* NoTertheleaa, a relatlTC atrength of the ocaq>lexea

formed by the rarloua iona under conaideration ahould be attainable from

theae more oonoentrated aalt aolutlona.

The data obtained from theae expérimenta at high ocnoentrationa

appear to be muoh more reliable than the data recorded in the fir at

expérimenta. In the low concentration expérimenta, the aalt aolutlona were kept very dilute in order to apply the Debye-HuTokel equation. Aa

a oonsequenoe, the increaae in oonoentrati on of mangane ae in the aolutlona,

due to the aalta added, ma hardly above the experimental error inrolved

in determining the total manganeae content in the aolutlona. Sowevor,

a definite Inoreaae in manganeae concentration ma obaerved in theae

expérimenta after adding the varioua aalta. 74

In the ezperlaente In whioh the salt oonoentrationa were muoh

higher, the inoreaae in manganeae oonoentration waa high enough to

eliminate a large part of the experimental error mentioned above. The

reaulta from theae expérimenta ahow that the F" iona and the HgPO^

iona are forming at ronger oomplexea with the manganoua iona than are the

NOg iona or the CIO^ iona. Furthermore, it appeara that the F io n a

are forming the atrongeat oomplex of the four, followed by the S^PO^

iona then the NOg iona and la at, the weakeat oomplex with C10^ iona.

(See Figure 16.)

Theae reaulta are baaed atriotly on the inoreaae in aolubility of manganoua oxalate obaerved after adding the varioua aniona to the

aolutlona. It muat be pointed out again that theae are only relative

oompariaona in oomplex atrength and no attempt was made to determine the actual stability oonatants for the ooaçlexea.

The oonplex uaed aa a referenoe point waa the EDTA oaaplex with the manganoua iona. Aa Indioated earlier, the stability constant for 11 t A this oomplex ia 10 whioh ia oonaidered a fairly atrong oomplex

for oomparative purposes here. In terms of inoreaaing oonoentration

of manganoua oxalate in solution, the muoh greater atability of this

ooa^lex ia quite evident in the plot of ^ o verses oonoentrati on of aalt in aolution (Figure 16). Here ^ o is the inoreaae in oonoentration of manganoua oxalate above that observed in a pure water solution. On a

oomparative basis it oan be stated that the aniona F , H.PO., NO and Z 4 5 ClO^ form relatively weak oomplexea with manganoua iona even at high

oonoentrationa of theae aniona.

From theae data about Mn** iona and other known faota on Mn*** iona it ia oonoluded that the complexing tendenoy of Mn^** iona far 75

•xoettd that of tho Mi** ion#, À ohook for oquillbrium in tfao salt solution was mad# on the last samples withdrawn in tho more oonooutratod salt solution oxporlmonts. Also a ohook for oquillbrium was mado on tho EDTA solutions*

The oonoontration of tho anions, CIO^, HO^, SO^, H^PO^ and F~, usod in tho rato-study oxporinonts ranged from 1,0M to 5.IM. Although the oonoontration of aniona in those oxporlmonts was higher than in the oomplex dotendCling experiments aboTO, the graph (Figure 16) from tho latter expo rimants indloatos Tory little oo^ lexing and no unusual behavior orsr the entire anion oonoentration range from 0»Q1M to 0.50M.

Therefore, it would not be unreasonable to assume that the Tory weak oomplexing tendenoy oarries over into the l.OM and 5,011 range.

The concentration of the EDTA solution was limited to approximately

0,6M due to low solubility of the EDIA salt* All data are tabulated in

T a b le s 7 and 8 .

Figure 17 shows a plot of log Kg for EDTA complexes of rarious metals verses number of electrons in atom. The log of Kg is given as

1., EDTA"*

Here, "M" is the metal under consideration. This graph shows the trend in stability of the EDTA oosplexes with metals In the +2 and *5 oxidation states. It is seen that the stability of the complexes with the +3 oxidation state of the metal is always muoh greater than with the *Z state. It smy be possible to carry these trends to complexesother 76 1.30

1.20

i.iO

1.00

0^0

OftO

0.40

0 ijO 2.0 40 70 Vjr X 10

Ftgurt 15. IncrooM in lotubiMty a t MnC,C^ in protpnct of vorioui tottt.

130 120

IK>

MX) #o #0

• - EDTA M m • - (MF •«TM 50 • -((Myo, M " O- NMOl Mm 40 * * Mm

30

20

0 004 O.W 0.20 024 0.32 CMMWtriWM «( %t/H

Figur* 16. IncrtOM in concontraTion of M on go n tw in aolution to •ott fffoctt. 77

T&bl# 7

Inorsas* in Solubility of In tho Prooonoo of Voriouo Salt#

,. §alV " S a lt Saaplo Cono. Cono. 10* Sumbor X 10®Vl M olan/U tar *Log 3* X 10*

N 1. 0 .0 0 2 .1 0 0.322 a 2 . 1.00 2 .1 0 0.321 96.9 C 5. 1.59 2.13 0.328 100.5 1 4. 2.08 2.14 0.381 108.4 O4 6 . 2.78 2.16 0.335 106.9 6 . 3.94 2.19 0.341 112.7

N 1. 0 .0 0 2 .1 0 0.322 a 2 . 1.00 2.09 0.320 96.7 H 5. 1.59 2.13 0.328 100.5 Og 4. 2.08 2.14 0.830 103.1 5. 2.78 2.18 0.338 107.1 6 . 3.94 2.16 0.335 1 1 2 .2

N 1. 0 .0 0 2 .1 0 0.322 2 . 0.56 2 .1 2 0.326 1 0 1 .1 S* 8 . 1.186 2.19 0.340 1 1 1 .0 O4 4. 1.410 2 .2 1 0.344 114.3 6 . 1.600 2 .2 1 0.344 116.7 6 . 2.084 2.25 0.352 123.5

1. 0 .0 0 2 .1 0 0.322 N 2 . 1 .0 0 2 .1 1 0.323 97.2 a 3. 1.59 2.14 0.331 1 0 0 .8 F 4. 2.08 2.15 0.333 103.5 5. 2.78 2.19 0.341 107.5 6 . 3.94 2.24 0.350 113.6

I 1. 0 .0 0 2 .1 0 0.322 a 2 . 1 .0 0 2.06 0.315 96.2 S2 3. 1.59 2 .1 0 0.328 1 0 0 .0 P 4. 2.08 2.13 0.328 102.9 O4 5. 2.78 2.15 0.353 106.7 6 . 3.94 2.18 0.339 1 1 2 .6

Log 8 * ■ Log S + 3.0000 uhoro S !■ tho oonoontration of Un ++ In molos por Utor. 78

T a b l# 8

laoreA ie in S o lu b ility of IhCgO^ In the Presenoe of Varioua Salta

S alt Mn Sample Cono. X 10^ Cono. X 10^ S a lt Ihtaber (m o lea/l.) (m o le a /l.) Log S* X 10 • e X 10®

V 1. 0 .0 0 2 .1 0 0.942 0.998 a 2 . 6 .0 0 2 .2 0 0.942 1.176 0 .0 0 C 5. 10.00 2 .2 0 0.942 1.971 0 .0 0 1 4. 60,00 2.40 0.580 2.441 0 .2 0 H 5. 100.0 2.72 0.495 9.990 0.52 6. 900.0 9.42 0.594 5.601 1.22

N 1 . 0 .0 0 2.24 0.950 0.947 a 2 . 5.00 2 .2 2 0.946 1.178 0 .0 0 H 5. 1 0.00 2.26 0.954 1.980 0 .0 2 4. 60.00 2.60 0.415 2.458 0.56 6 . 100.0 2.92 0.466 9.942 0 .6 8 6 . 900.0 9.80 0.580 5.614 1.56

M 1. 0 .0 0 2 .2 0 0.942 0.998 a 2 . 5.00 2 .2 0 0.542 1.176 0 .0 0 Hg 5. 10.00 2 ,2 0 0.542 1.571 0 .0 0 P 4. 50.00 2.76 0.441 2.471 0.56 O4 5. 100.0 9.66 0.564 9.586 1.46 6. 900.0 6 .6 6 0.824 5.716 4.46

1. 0 .0 0 2.24 0.950 0.947 N 2 . 6 .0 0 2.26 0.554 1.185 0 .0 2 a 5. 10.00 2.94 0.969 1.991 0 .1 0 F 4. 50.00 2.90 0.462 2.482 0 .6 6 5. 100.0 9.82 0.582 9.395 1.58 6. 900.0 6.94 0.841 5.726 4.70

1. 0 .0 0 2 .2 0 0.342 0.998 E 2 . 10.00 1 1 ,0 1.041 9.74 8.80 D 9. 50.00 40,9 1.612 8.12 58.70 T 4. 100.0 72.0 1.857 11.90 69.80 A S. 900.0 122.0 2.065 18.70 119.80

o " 0 - o.o where o la the obaerved oonoentration of Mn** In a a lt aolution and la the oonoentration of Mn** in pure water. 28

24

0 3 + 2 Sfot* in O.IN KCI O - +3 in 0.1 N KMOj

20

28 3 0 32 5 4 38 3 8 40 No. Eleefrene h Atom Figurt 17. Cofnporison Of stability constants of complexes of metal ions. 80 than EDTA conplezaa* In any avant, the graph ahowa that the +3 oxidation state of the metals is highly favored in oomplex formation*

Separation of Manganeae Complexes Using Domex 1-X6 Resin

The oonclusxon reached in the previous section was that the +2 state of manganese forms relatively eeak complexes with the anions P",

HgPO^, NOj, CIO^ and SO^. The next step was to carry out an inves­ tigation to determine whether the complexes formed by the higher oxidation states of manganese could be separated using Dowex 1-18 anion exohange resin.

Discussion of Results Obtained from the Resin Runs

Before disoussing the final results obtained from these runs, it is desirable to first describe conditions that are believed to exist in the solutions used in the experiments. Calculations of results will then be made based entirely on these conditions and the concentrations of the various species put into the solutions. These results will be compared with the actual experimental results*

It was observed, from speotrophotometric studies made on solutions containing manganous ions and permangenate ions in hydrofl’-ori c acid, that a very rapid decomposition of permanganate ions occurred in these solutions immediately upon adding the potassium permanganate. A natural assumption that would account for this observation would be that the fluoride Ions were complexing with some intermediate oxidation state of manganese in this system and hence shifting the equilibrium or at least accelerating the rate in a direction by the mass action effect*

Anparently, in order to cause such a rapid decomposition, the complex being formed must be a verj' stable one. The great stability of the 61

ocnqplaz i i mor# apparent when i t ia noted t h a t , aven a f te r g ettin g

aereral week#, the aolutlona oontaining the fluoride iona had no

Tiaihle precipitate of manganeae dioxide in them* In all other aoid

aolutlona (except phoaphorio aoid aolutlona) of identical potaaaium permanganate and manganoua a u lfa te oonoentrationa, th e appearanoe o f a preoipitate of manganeae dioxide waa almoat alwaya obaerved within te n houra.

Thua, in view of the probable presence of an apparently atrong

complex in the hydrofluoric aoid aolutlona, it waa thought that this

complex of manganeae oould be separated from the *2 oxidation state of smnganeae by using an anion reain. In the event that the complex oould be isolated using an anion resin, it would then seem that the actual

oxidation state of manganese in the oomplex oould be determined by observing the activities in the different fractions* Therefore, the problem waa to isolate the oomplex. To aoooaipliah this, Dowex 1-X8 anion exohange reain waa uaed.

The so lu tio n s were prepared In th e follow ing manner. The manganoua aulfate solutions, both aotive and inactive, were put in the sample flask first, then the hydrofluoric aoid added, and then the potassium permanganate. This «as the general prooedure for nuiking up the samples, th a t is , the manganoua iona were added f i r s t , followed by the aoid and then the permanganate. The order of addition was very important in thia work. The only exception to the above order of addition was in the run in whioh the Mn -MnFg isotopio exohange was detersiined. (The asterisk BZ

represent* the Initial radioaotive apeoies. Aotuallj, the Nn** ion probably formed a weak oomplex in the 5.IM aoid eolation, but in this disou#sion the Mn*^ ions will be represented as unoomplexed,) The

Mn**-MnFg exohange will be disoussed later in this seotion.

The solution used in the ion exohaage work always had an excess of manganous sulfate present. For the first run made, the oonoentration

of manganous sulfate was 8.42 I the potassium permanganate

oonoentration was 1.00 X 10”®M and the hydrofluoric acid concentration was S.lOH.

In the presence of the large excess of manganous sulfate (8.42 I

1 0 as compared to potassium permanganate ( 1 .0 0 % I0 "^lî) » the l a t t e r

always deoomposed cmaplately in the 3.lOK hydrofluoric aoid solutions.

Therefore, eluent from the resin column was clear. There should be no

reaction with the resin. The volume of the samples taken from the column was 0.20 ml. These samples were counted for activity and a plot was made of the activity of the sample verses the total number of m illiliters of sample withdrawn. As seen from the plots in Figures 16 and 19, a separation was attained in each case where hydrofluoric aoid was used.

In the first run. Figure 18 (a), there was a considerable overlap

of the peaks. This was noted before the other runs were made. There­ fore, in order to get a better separation of the peaks, the solutions

for the other three runs were put through the column at a slower rate than the first solution. The resulting separations were much better as indicated by the plots in Figure 16 (b) and (o) and Figure 19, 8S

To h« oertain that th# +2 oxidation state of nanganese was not being oxidised or reduced lay the resin, several runs were made in whioh only Un ions in 5.10V. hydrofluoric aoid were eluted from the column.

In each case, a single well-defined peak was obtained. The recovery of manganese was 100% w ithin experim ental e rro r. Figure 19 (ej shows a typical plot obtained from the data. Thus, it appears that the resin had no e ffe c t on the manganous ions in so lu tio n .

Having separated the oomplex KnFg from the manganous lone, the next problem was to determine whether the oxidation state of the manganese loos in the costplex could be determined from the data. To accomplish this, the following procedure was used.

Three experiments were carried out using various concentrations of manganous sulfate and potaselum permanganate. The hydrofluoric aoid concentration was kept at 5.10V in all experiments. The first solution to be eluted from the resin column was initially 8.42 I lO'^V in manganous sulfate and 1.00 X lO'^V In potassium permanganate; the second solution was initially 8.42 X 10~^U in manganous sulfate and 1*05 % lO'^V in potassium permanganatej the final solution was initially 7.02 X lO'^M in manganous su lfa te and 1.00 X 10*^K in potassium permanganate.

The values of these concentrations are very Important for purposes of calculations.

It was first necessary to consider some possible reactions taking in the solutions and complexes formed from these reactions. Two possible overall reactions would be 84

(i a ) 4Mn** + MnO^ =. 5Mn*** (ocm plex)

or

(IIA ) 5Mn** + 2 1 1 nO^ — 6 kn** (complex)

In the flrct reaction, equation (IA), 4 Mn^ react for every one UnO^ ion. Aaauming the reaction goes to ooaipletlon and alow exchange between

Mn*^ and MoFg, then the activity obaerved in each fraction (i.e., the

Mn** fraction and the MnFg fraction) ahould be

, , , , V (Mn**) - 4(Mn0 4 ) (IB) (Aot)jjjj++ ■ (Act)^ ------and

(Aot)j(uf- - (Aot)^ - (Act)j^++

The parentheaia around the iona indicate oonoentrationa; (Act)^ ia the total act :vity put in the aolution. Uaing the oonoentrationa given above fo r the aolutlona and the oorreaponding total activity observed for each

aolution, the calculated results would be (experiment 1 shows sample oalculati on)

8.42 10'® - 4.00 10"* Exp. 1 (Act) * 55,800 ------» 17,480 cpm Mn 8.42 10"*

(A ot)^E • 33,500 - 17,480 » 18,820 opm 6

Exp. 2 (Aot)^^++ = 17,200 opm * 16,965 opm 6

Exp. 5 (Aot)jjQfV 16,120 cpm (Act) » « 21,560 opm 85

A comparison of the above calculated results with the experimental results Is given in the table below*

Table 9 i Data from Separation of Mn and MnF^ Using Dowex 1-X8 Resin

Experiment Calcd'd Activity Calo'd Activity Exp. Observed Number Usine Equation (IA) Using, Equation (IIA) Activity

Mn** Frac. MnF* Frac. Mn** MnF* Mn** llnFg

1 . 17,480 15,820 27,370 6,930 19,136 14,170

2 * 17,200 16,965 27,760 6,416 16,876 17,290

3. 16,120 21.360 29,470 8 ,0 1 0 16,726 20,766

Comparing the calculated results witb the experimental results shows quite obviously that the experimental data correspond much closer to the reaction occurring in equation (IA) than in equation (IIA)* In fact, within experimental error, the data show that the reaction is almost oompletely as represented in equation (lA)* The larger error observed in experiment 1 can be assumed due to Incomplète separation of the two fractions as indicated by the overlap of peaks in Figure 18 (a)*

In spite of this error the results agree much closer to the results calculated from equation (lA) than (iiA)• Therefore, the conclusion reached is that the oxidation state of manganese in the oomplex is the

+5 oxidation state.

The results obtained from the experiments above are in general agreement with the observed stability of the +3 and +4 oxidation states of manganese in solution* The +3 state is unstable in solution 86 unless stsblllted by ocnplexingj the +4 state Is generally alimys unstable in solution, preferring to preoipitate as sianganese dioxide*

(See lotrcduotion*)•

Experiments mere also made using perohlorio aoid instead of taydrofluori 0 aoid, but no complex oould be separated from these

solutions* (See Figure 19 (b).) This would seem to indioate that the

CIO^ ion does not form a stable complex with the *3 oxidation state of manganese* Most likely a weak oomplex is formed by the CIO^ ions in the ++ » solutions, otherwise, the rate of decomposition of the Mn -MnO^

system observed by speotrophotometric methods would be even more difficult to explain. It would seem, if the +3 oxidation state were not oomplexed at all, that the appearanoe of a preoipitate of manganese dioxide would be much quicker than observed, sinoe in pure water (no acid) the preoipitate appears almost immediately.

I t was reri-ised th a t the outcome of the problem above, from the viewpoint of observed sastple activities, depended entirely upon the ++ ; rate of isotopio exchange between Mn and MnFg. A very fast exohange

or even a complete exohange within 15 minutes or so, would give incorrect results* In view of the importance of this exchange, the following experiment was made in order to determine the actual rate of exohange.

Manganous su lfa te was f i r s t added to the fla s k , then hydrofluoric aoid followed by potassium permanganate. The color of the potassium permanganate disappeared immediately. From the results of the previous experiments, the solution contained manganese ions in the +2 and +5 87 oxidation state* *dth the latter oxidation etate in the form of a stable oomplex. The solution was allowed to set for approximately 50 minutes to assure equilibrium conditions. Then carrier-free radioactive manganous ion* were added to the solution. The final oonoentrationa of the vari ous speoies present in the solution weret (HnSO^) * 8*42 I l O '^ l (KMnO^) - 1.00 I and (HF) - 3 .ION.

After setting, with occasional stirring, for one hour and ten minutes, the solution was passed through the reain column. The results are tabulated below. Figure 19 (a) shows the peaks obtained from the sample activities taken from the column.

Table 10

Data on the L a b ility of Mn^^ and MnP^

Experimentally Observed Activity Caloulated Act. for Cosqjlete Exch*| Mn^^ F raction Fraction Mn** F ractio n MnFg F raction

27,599 opm 10 ,,061 opm 17,600 opm 19,900 cpm

Based on the +3 oxidation state being present in the stable oom­ plex, the activity that should have been observed in each fraction was calculated as follows

(izB)

8.42 I 10*^ - 4.00 X 10 (Act)^++ - 37,600 — ■ " J, " " 17.600 cpm 8.42 X K)'® + 1.00 X 10"®

(Aot)jjjjp* * 37,600 - 17,600 - 19,900 cpm 88

It warn quite upperent from the graph and from the comparison of the actlTities of each fraction in Table IQ that complete exchange did not occur* The obaerved activity of the Mn** fraotlon was muoh less than expected for a complete exchange* From the experimental activity values, ** * the half-time for the Mn -UnF isotopic exchange «as found to be 0 approximately TO minutes*

It should be noted hero that the rate observed over an hour and ten minute period, in the experiment above, should be a maximum rather than a minimum rate for the Mn**-MnT. exchange, that is, the counts per 0 minute observed in the MnF* fraction should be a maximum value* The o reason for this is apparent in the experimental conditions* The radio­ active manganese added to the solution was in the *2 oxidation state, and the manganese was carrier-free. The two most important conditions that

■dght have existed In the solution, which would have caused errommous results, would be, first, if the solution had not yet attained an equilib­ rium between Mn**^ and MnFg, and second, i f the re a c tio n of suinganous a sulfate with potassium permanganate to form MnFg was not completed*

But both of these conditions favor higher activity in the MnFg fraction + + I than would have been observed in a true Mn -MnFg exchange* In view of the results obtained from the previous experiments in this section, there m would be no reason to believe that less of the MnFg complex was present in this experiment than in the others since the composition of the solu- 4-f ■ tion in the Mn -MnFg exohange experiment was identical to two of the solutions used in previous experiments* Also these previous experiments support the conclusion that complete exchange does not occur instantaneous- ++ a ly between Mn and MnFg since none of the experiments gave results which would correspond to a complete exchange. 69

M th the eitftbllihment of aa Incomplète exohange between Wn** and MnF* and of a half-time exohaage of at least one hour, it follcws that the error ooourrlng in experiments 1, 2 , and 3 due to th is exohange should be smell, that is, the oontribution to the aotlvlty values of each fraction in these experiments from the Mn*^-MnFg exchange should be small. Before continuing, it is important to distinguish between the two reactions that oould be occurring in the solutions in experi­ ments 1, 2, and 3. As explained earlier, the manganous ions, both active and inactive, and the hydrofluoric add were put into the flask first. The potassium permanganate was added which reacted with the s manganous ions to form the MnFg complex. The oomplex contained a considerable amount of activity because of the radioactive Mn** ions in the solution prior to adding the potassium permanganate. However, the activity In the oomplex, due to the reaction of permanganate iona with manganoua ions, should not be as great as that caloulated for an instantaneous and complete exchange between Mn** and MnFg ions- (See equations (IB) and (lIB) above.) Therefore, a second reaction in­ volving an exclmnge of radioactive ions between Mn** and MnF* would have been occurring also. But this exchange should be quite slow because of the large amount of activity already present in the MnF* 6 oomplex. Thus, the activity contributed to the MnF^ complex from this exchange should be a small percentage of the total activity.

The deviations of the observed activities from the calculated values (see Table 9) could be due to this exchange mentioned above.

In summary, the follow ing conclusions can be made from the experimental results: 90

1. The main Intermediate oxidation state of manganese present In the Mn -UnO^ -HF system Is the +3 oxidation state.

2. The exohange between Mn^^ and MnFg is not instantaneous but was found to have a half-time exchange of approximately one hour. 5 0 0 0

4 5 0 0 Soin 0 i56 nM 0054 V MnSO, Soin 0 (3 0 ml MnSO, 0 470 ml VnSQu CF 0 4 7 0 ml MnSO, 0 4 7 0 ml MnSO, 4 0 0 0 0 355 ml HF 0 355 ml HF 0 020 05C' M KMnO, 0021 mi kMnO, 0 0 2 0 ml KMnQ,

350 0

5 0 0 0

2000

5 0 0

OOO

5 0 0

C * 1» 0 20 30 4 0 6 0 70 30 4 0 50 6020 3 0 40 5 0 60 70 ml Solution ml Solution mi Solution ID Figure 18 (o),(b) and (c) Spectro obiomed from eluent token from onion exchange column 500 0

‘ - W"SC, 3 Î55 "1 HF 34T - V*50, HF SofxJord i2,6i6 CPM 4 0 0 0 C ?00 mi X MnQ, 0 Î5C Gout's *rom resm 2,790 CPM CdTC ml WnSO, 0 ?50 m xMnC

2 2500 o

2000

1500

1000

500

2 0 30 5 0 6 0 70 30 5 0 6 0 20 30 5 0 60 70

Figure S (o),{b)and (c) Spectra obtained from eluent token from onion exchange column. IP ro 95

SPECTROPHOTOMETFTC STUDIK5 OF THE SYSTa* Mn**-KnO^

A «tudy was made of the rate of deoomposition of permanganate in the pr.esenoe of nanganous ions and several different aoids. The concen­ tration of the two manganese species and the acids were varied over a fairly wide range In order to obtain a more complete understanding of the processes occurring in the solution. The rate of decomposition of permanganate in perchloric acid solutions was studied thoroughly. A

less thorough study was made In the other acid solutions. The acids used were : perchloric acid, nitric acid, sulfuric acid, phosphoric acid and hydrofluoric so ld .

Experiments were also made in acetic add and in mono and tri- chlorioacetic acids, but in all of these solutions the rate of deoom- positj on was too fast to follow. The acetic acid solutions imnediately started turning a clear deep wine color. Witiiin lb minutes after mixing, the solution had turned a muddy brown color with the formation of a precipitate clearly visible. The same general decomposition procedure was observed in the mono-chloroaoetic acid solutions, except that the deep wine color did not appear as quickly whereas a precipitate appeared more quickly (within ten minutes after mixing). Possibly the presence of the or lord ne atom in the mono-chloroacetic acid caused the permanganate to decompose more rapidly in this solution than in the acetic sold solutions, however, the increase in the hydrogen ion oon- oantrat: on in the mono-chloroacetic acid solutions over that in the acetic acid solutions apparently is not great enough to delay the appearance of the precipitate. It would appear that the acetate ions are forming a slightly stronger complex than the mono-chloroacetate ions. 94

Tri-ohloro«cetio acid is a atrong acid* Ko precipitate appeared

In these solutions for at least S5 minutes after mixing* However, the solutions were almost completely decomposed within 15 minutes after mixing as Indicated by the nearly colorless solutions. It appears that the more chlorine atoms attached tc> the acetjc acid the more rapidly the permanganate decomposes. Permanganate ion is reduced by chloride ion so possibly the same reducing ability is present when the chloride ions are attached to the acetate ions. The reason for the delayed appearance of the precipitate could be due to the much higher hydrogen ion concentration in the trichloroacetic acid solutions.

As a final step in the attempt to solve this problem, the rate of decomposition of pemanganate in the Mn**-I,inO^-acid system was studied in the presence of the sodium salts of the acids used in the experiments.

By adjusting the pK to the proper value, the effect of the anions of the various acids on the rate of decomposition of permanganate was determi ned.

All data is tabulated at the end of this discussion. Graphs obtained from the dota are given in Figures 20 and 21 at the end of this section.

Discussion of Results Obtained from Spectrophotometric Studies

The spectrophotometric studies in perchloric acid solutions were used for determining some of the parameters that effect the kinetios of the system Kn*^-Kn0^-EC10^, One reason for using the experiments in perchloric acid solution was that the perchlorate ions have the 95 least oomplexlng tendencies of any anions studied with the manganese ions in solution. Therefore, the various oxidation states of manganese were more free to enter Into reaction with eaoh other, resulting in minimum oomplexlng effects on the rate of reaction.

The first attempt to determine the order of réaction with ++ * respect to Mn and MnO^ was made by taking the slope of the curves at time t ■ o from the optical density verses time plots. In one set of curves the initial UnO^ concentration was kept constant adiile the initial Mn^ concentration was varied. In the second set of curves the initial MnO% concentration was varied. In all runs, the acid concentration vas very high compared to the other species present and therefore, the hydrogen ion concentration was assumed to be constant.

Because of the high concentration of acid, the ionic strength of the solutions was also assumed to remain constant and thus the activity coefficients of the various species remained constant. The specific rate equation isi

By t.lüng th. slop. - ~ »t tlm. t * o, th. Initial conoentration of at Mn*^ and MnO^ could be used. Since in one set of curves the initial

UnO^ concentration was constant and in the other set of curves the initial Mn** concentration was constant, a plot of log (- ■“ ) verses log of initial concentration (of the species) which was varied for eaoh set of curves should give a straight line with a slope equal to "a" for oonstant UnO^ concentration. A near straight line was observed in 96 the Ifctter oaso. The value of "h" was approximately 0.85. However, in the former case the value of "a" varied from about 1 .2 to 1 .6 in the higher oonoentrated and exoess Vn** ion solutions to about 0.67 in the more dilute solutions in which the ICnO^ concentration was equal to or greater than the Mn** concentration.

The procedure for determining "a” and "b” was net very accurate e sp e c ia lly in th e more concentrated so lu tio n s as was indicated by a greater scattering of the points in those solutions. However, a definite trend to higher values of ”a" was observed in the graph (Figure 22).

Certain experiments showed some interesting results. For example in the experiments in Wiich the Initial Mn** concentration was 2-l/2 tistes the initial MnO^ concentration it was found that the reaction followed second-order kinetios, i.e., a*b*l. A plot of — verses time C ' ' I for the experis'ents gave straight lines with average values of k equal to 0.^7 in l.OM perchloric acid, 0.72 in E.OM acid and 1.93 in 3. IK acid. A straight line was also obtained in the plot of verses time for the solutions in which the Mn** concentration was twice the

MnOj concentration (all concentrations given in moles per liter).

The value of "c" was calculated from the opticsl density equation d = k c l.

No straight lines were obtained in the plot of data from the other experiments where the permanganate concentration exceeded the manganous ion concentration. Apparently the assumption that a=b*l for these solutions i 8 no longer valid. (It was stated earlier that "a" was 97 approximately 0.67 in these solutions. This seems to be verified in the deviation of the plots from straight lines in these solutions.)

Another reaction most likely predominates in these solutions. No further work was done on this parti oular area of study.

The next area of experimental study was concerned with the rate of deoompositioD of permanganate in the presence of various acids.

The spectrophotometric data Indicated that the rate of decomposition was approximately the same in perchloric and nitric acids. The rate was slightly faster in sulfurio acid and much faster In phosphoric acid and the most rapid in hydrofluoric acid. A typical comparison of rates is shown in Figure 2 0 . This was the general observation in all such decomposition curves. These décomposition rates can be easily correlated with the complexIng ability of the an one as determined from the com- plexlng studies carried out in this researci.

It was observed in the study of oomplexlng ability of the aniens - _ ■ CIO^, SO^, H2P0^ and !■ with manga nous Ions th a t the order of increealng complex stability was in the order given above. This order also carries over into the manganic ions (Mn®* as inoioated in Figure 17 on page 92* The stability constants are also greater for the manganic ions than for the manganous ions. Since the manganic ions were found to be the main intermediate oxidation state in the system Mn**-MnO^-acid, the increasing stability of the complexes of the various anions with the +3 state resulted in a corresponding increase in the decomposition rate of the permanganate ions. 98

To verify the above results several experimente mere made to determine the rate of décomposition of pennanganate ion in the presence of the sodium salts of the anions dlsoussed in the previous paragraph.

The pH of the solutions were adjusted, using a glass eleotrode, so that the hydrogen ion concentration was approximately the same for eaoh experiment. Thus, the relative effect of the anions on the decomposition rate could be determined. It was found that the perchlorate and nitrate

Ions haa no noticeable effect on the rate. The sulfate ions had some effect. The di-hydrogen phosphate lone and the fluoride ions had a much greater effect on the rate. These results agree with the overall results obtained in the acid systems discussed on the previous page.

This completes the spectrophotometric studies on the system

Mn**“î£nC”- acid and on the acid salts of this system. Many experiments were made in these various acid systems. A complete tabulation of data is given in the following pages. 99

T&bl# 11

Composition of Solutions for Exp#rim#at# I through XL

MnO“ Mn** Ifci** Mn** Mn** Mn** Mn** Exp* Aoid Cone** Cone. Cone. Cone. Cone. Cone. Cone. Number Aoid Normality X 104 X 104 X 10* * 10* X 10* X 10* X 10* OD-1 CD-2 OD-3 CD-4 ÔO—5 QD-6 I HC104 3.1 10.0 20.0 10.0 5.00 2.00 1.00 11 H 3.1 8.03 10.00 5.00 2.00 1.00 III •t 8.1 5.00 12.5 10.0 5.00 2.00 1.00 0.500 IV n 3.1 3.50 17.5 8.75 5.00 3.50 2.00 1.00 V * 3.1 2 .00 5.00 2.00 1.00 0.50 VI n 2.0 5.00 12.5 5.00 2.00 VII n 2.0 3.50 17.5 8.75 3.50 2.00 VIII n 2.0 2.00 5.00 2.00 IX m 1.0 5.00 12.5 6.00 2.00 X m 1.0 5.50 8.75 3.50 2 .0 0 JJ ft 1.0 2 .0 0 5.00 2 .0 0 1.00 XII E3P04 3.1 5.00 12.5 10.0 6.00 - n n tt 3.1 3.40 17.5 8.75 5.50 XIV « 2.0 6.00 12.5 10.0 5.00 XV « 2.0 3.40 17.5 8.76 8.50 XVI M 1.0 6.04 12.7 10.0 5.00 XVII n 1.0 3.45 6.91 3.54 1.47 XVIII # 0.10 6.00 6.00 2.50 1.00 XIX H 0.10 3.40 3.50 1.00 0.50 XX H2S04 3.1 5.00 12.5^ 10.0 5.00 XXI ft 5.1 5.40 17.5 8.75 5.50 XXII ft 2.0 6.00 12.6 10.0 5.00 XXIII * 2.0 8.60 17.5 8.67 3.54 x n v ft 1.0 5.00 12.5 10.0 6.00 XXV ft 1.0 3.40 8.76 3.50 1.80 XXVI HlOj 5.1 6.00 12.5 10.0 5.00 XXVII m 5.1 5.40 17.5 8.75 3.50 XXVIII n 2.0 5.00 12.5 10.0 5.00 XXIX tt 2.0 5.40 17.6 8.75 S.SO XXX tt 1.0 6.00 12.5 10.0 6.00 XXXI tt 1.0 5.40 8.76 8.50 1.50 XXXII HP 3.1 5.00 12.6 6.00 XXXIII tt 2.0 5.00 12.5 5.00 x x n v tt 1.0 5.00 12.5 5.00 XXXV* equiv* to 2 .0 H 3.50 17.5 3.50 (in HCIO4 ) XXXVI tt tt tt 3.50 17.5 3.50 (in HHO5) XXXVII tt " l.OH 3.50 17.6 3.50 (in H2SO4 ) x g g i i i " " 2.ON 5.50 17.5 3.50 ( in HNOJ XXXIX H5P04'' H2PO4 - 6.00 12.6 5.00 12.5 5.00 12.6 5.00 XL HC104 5.1 6 .00 12.5 12.5 12.6 12.5 12.5 12.5 XXXIVA HP 3.1 (3.66-3.51) 8.90 3.62 * Experiment# XXXV throng XL maide In #alt# of aoids on right, **In #aoh #xperim#ntj th# KlfnOg oono* remained constant (exoept in experiment XXXIVA) srhlle the Mn** ooncentration was varied as Indioated. T able 12 Optical Density (OD) of Potassium Permanganate in Hater and Various Àoids

Potassium Permanganate Cono. 5.00 X Molar X iO-4 Molar Experimental Conditions cftJ-I T OD-2 ¥ " w - y T o w OD-5 I Cb-6

Mo aoid Just pure water 1.250 1.250 1.260 0 .8 6 5 0 .8 6 5 5 .1 M BCIO4 50 1 .175 65 1.172 125 1.172 50 0.805 66 0 .8 0 0 125 0 .796 2 .0 M HCIO4 SO 1.185 65 1.185 125 1.185 50 0.812 65 0.810 126 0.800 1 .0 H HCIO4 so 1 . 2 0 0 65 1 . 2 0 0 125 1 . 2 0 0 50 0 .8 2 0 6 6 0 .8 2 0 125 0 .8 1 5

5 .1 H HMO3 65 1 .225 1 0 0 1.226 160 1 .2 2 6 66 0 .8 5 7 100 0 .8 5 7 150 0.858 2 . 0 N HMOs 65 1.220 100 1.220 150 1 . 2 2 0 65 0 .8 7 6 * 100 0 .8 7 7 * ISO 0.871* 1 .0 H HNO5 65 1 . 2 2 0 1 0 0 1 . 2 2 0 150 1 . 2 2 0 65 0 .8 5 5 xoo 0 .8 5 1 ISO 0 .851

5 .1 N 5 2 &O4 80 1.226 1 2 0 1.226 166 1 .2 2 5 80 0 .8 2 5 120 0 .8 2 0 165 0.816 2 .0 M M2 MO4 80 1.224 120 1.226 165 1 .2 2 5 80 0 .6 5 9 ISO 0.858 166 0.838 1 . 0 M H2 8 O4 80 1.224 120 1.226 166 1.225 80 0 .8 5 6 120 0 .8 5 6 165 0 .8 5 5

5 .1 N H5 PO4 6 1.220 30 1.220 60 1.220 5 0 .8 2 8 so 0 .8 2 6 60 0 .8 2 8 2 .0 M H5 PO4 6 1.220 30 1.220 60 1.220 5 0 .8 3 5 30 0 .8 5 5 60 0 .8 3 2 1 .0 H H5 PO4 5 1.220 50 1 .2 1 5 60 1 .2 1 5 5 0 .8 2 8 50 0.825 60 0.836 0 .1 0 N H3 PO4 6 1.250 SO 1.225 5 0 .8 5 0 50 0 .8 2 8

2 . 0 N HF - 1.225 ISO 1.185

* See footnote on follosdng page. % Values high due to higher concentration of potassium permanganate in these solutions. T able 13 Optical Density of KMnO^ as a Function of Time in the System Wn**-MnO^-Aold**

Exp. No. T* OD-1 T OD-2 T CD-3 T 0D'*4 T CD-5

40 1.78 49 1.99 54 2.08 64 2.09 52 2.08 49 1.74 59 1.97 64 2.05 64 2.08 62 2.08 59 1.64 71 1.92 78 1.99 79 2.06 77 2.07 I 70 1.54 87 1.89 115 1 ,^ 6 113 2.05 112 2.08 86 1.68 97 1.86 145 1.95 148 2.02 149 2.07 97 2.02 119 1.78 178 1.90 178 2 .0 2 177 2.05 107 2.09 166 1.86 201 1.89 201 2 .01 199 2.06 183 2.00 281 1.85 281 1.96 279 2 .0 1 206 2.08 326 1.83 526 1.93 322 2 .0 1 373 1.81 579 1.85 376 1.96 429 1.75 429 1.80 427 1.87 456 1.69 460 1.78 456 1.84

15 1.77 13 1.77 19 1.79 20 1.74 29 1.70 31 1.76 38 1.76 38 1.72 45 1.67 45 1.73 51 1.75 54 1.68 II 60 1.62 00 1.71 63 1.75 66 1.70 76 1.60 75 1.69 108 1.73 108 1.69 113 1.55 113 1.67 142 1.70 144 1.66 160 1.42 161 1.64 163 1.68 183 1.65 194 1.46 194 1.57 223 1.67 225 1.62 235 1.45 236 1.53 284 1.65 285 1.59 283 1.23 263 1.48 535 1.57 336 1.51 541 1,68 371 1.54 374 1.47

♦T is the time 1 (in minutes/ a fte r mix! nt: the aolut I on ; D is the optical density for solutions 1, 2 , 5, e tc . Conmoaition of solutions given in Table 11. • ♦ Acids used in the experiments were HCIO 4 . HNO3 . H5PO4 and HF. Table 13 (Contd.)

Exp. No. T OD-1 T 00-2 T OD—3 T OD-4 T OD—6 T OD—6

14 1.11 25 1.12 26 1.16 26 1.15 24 1.16 79 1.14 20 1.09 58 1.09 47 1.40 48 1.14 49 1.16 118 1.13 32 1.06 56 1.06 86 1.10 87 1.13 86 1.16 203 1.12 47 1.02 66 1.05 117 1.08 119 1.12 118 1.14 256 1 .12 65 0.97 75 1.01 146 1.06 146 1.11 146 1.13 308 1.10 96 0.89 99 0.96 203 1.02 204 1.08 203 1.11 370 1,09 ITT 128 0.81 114 0.94 255 0.98 255 1.07 264 1.10 465 1.08 146 0.82 144 0.89 306 0.94 308 1.06 307 1.08 574 1.06 150 0.83 185 0.84 368 0.91 368 1.02 368 1.07 727 0.81 169 0.90 228 0.87 464 0.91 466 0.98 464 1.04 194 0.89 286 1.04 674 0,92 675 0.92 674 0.99 232 0,75 734 0.48 754 0 ,6 8 754 0.75

16 0,75 11 0.80 11 0.81 10 0.81 20 0.80 21 0.81 29 0.69 21 0.78 21 0.80 20 0.80 88 0.77 87 0 .0 0 42 0.65 52 0.77 32 0.79 52 0.79 157 0.75 135 0.78 61 0.60 44 0.76 48 0,78 49 0.78 198 0.75 196 0.78 91 0,52 56 0.75 66 0.77 56 0.78 261 0.72 250 0.76 IV 124 0,49 72 0.71 75 0,76 72 0.77 312 0.70 311 0.75 141 0.48 92 0 .6 8 92 0.73 93 0.75 372 0.69 571 0.74 145 0.48 115 0 .6 6 116 0.72 114 0.74 456 0.67 454 0.75 165 0.46 133 0,64 133 0.71 132 0.72 529 0.64 628 0.71 190 0.42 148 0.62 148 0.70 147 0,71 228 0.55 209 0.67 209 0.67 208 0.69 251 0.54 251 0 .6 6 251 0 .6 8 270 0.64 270 0.64 269 0.67 300 0,53 300 0,63 299 0 . 66 332 0.64 333 0.61 532 0.65 424 0.51 426 0.69 426 0.63

g INJ K>S

TabU 15 (Cootd.)

Exp. No. T OD-1 T CD-2 T OD-3 r OD-4

16 0.45 14 0.46 13 0.46 14 0.50 31 0.44 30 0.44 28 0.45 28 0.46 59 0.42 56 0.43 58 0.44 58 0.45 V 103 0.40 102 0.41 101 0.43 106 0.44 160 0.37 169 0.40 157 0.42 151 0.43 216 0.35 216 0.39 217 0.41 210 0.43 306 0.32 305 0.37 303 0.39 SCO 0.41 414 0.29 413 0.56 413 0.38 406 0.40 619 0.25 618 0.32 617 0.35 617 0.38

61 1.12 50 1.16 48 1.16 100 1.05 98 1.15 97 1.16 168 0.98 167 l . l l 165 1.14 VI 201 0.94 231 1.08 230 1.13 252 0.96 287 1.06 286 1 .1 1 260 1.01 380 1.04 379 1.09 289 0.98 527 0.98 426 1.07 580 0.76 684 0.96 583 1.07

17 0.80 60 0.80 15 0.82 59 0.82 31 0.77 113 0.77 94 0.76 113 0.80 44 0.76 178 0.73 153 0.73 178 0.79 65 0.73 242 0.69 213 0.70 242 0.78 94 0 .6 8 298 0 .6 6 275 0.67 298 0.76 VIT 126 0.63 390 0 .6 6 322 0.64 392 0.75 143 0.61 536 0.61 353 0.62 536 0.73 146 0.61 586 0.61 590 0.73 166 0.62 426 0.62 191 0,64 450 0.67 230 0.65 490 0.-72

13 0.47 12 0.47 59* 0.82* 91 0.46 91 0.45 112 0.80 ISI 0.46 150 0.45 179 0.80 211 0.45 211 0.43 241 0.79 273 0.44 273 0.41 296 0.79 vin 322 0.44 320 0 .4 0 392 0.78 351 0.43 350 0 .4 0 536 0.76 384 0.43 585 0 .9 8 589 0.76 423 0.43 422 0.39 449 0.43 449 0.38 468 0.42 488 0.38

♦These oolumne belong in Experiment V II. Composition of th is solution given in Experiment vii in the Composition Table 11. 104 Table IS (Contd.)

Exp. No. T OD-1 T OD-2 T OD-3

28 1.14 29 1.17 28 1.18 48 1.14 48 1.16 48 1.17 82 l . l l 81 1.15 82 1.16 142 1.07 142 1.13 143 1.14 177 1.09 196 1.13 197 1.15 IX 197 1.14 262 1.13 262 1.16 267 1 .1 1 286 1.15 286 1.14 286 1.02 314 1 .1 2 316 1.14 316 0.92 348 1.13 349 1.14 349 0.80 385 1.17 387 1.13 585 0.70

14 0.83 15 0.84 15 0.84 54 0.85 34 0.83 35 0.64 67 0.81 67 0.85 67 0.83 90 0.81 90 0.82 90 0.83 115 0.80 115 0.82 115 0.82 154 0.80 154 0.81 164 0.82 182 0.83 182 0.81 182 0.81 238 0.83 237 0.80 237 0.81 272 0.87 271 0.80 271 0.81 300 0 .8 6 299 0.80 500 0.81 534 0.81 533 0.81 334 0.81 373 0.76 372 0.82 373 0.80

26 0.48 26 0.47 24 0,47 67 0.47 66 0.47 65 0.47 102 0.47 101 0.46 100 0.47 XI 169 0.46 159 0.46 157 0.46 219 0.45 217 0.45 215 0.46 283 0.44 282 0.46 281 0.46 343 0.44 543 0.46 341 0.46 496 0.43 497 0.45 496 0.45

16 0.57 14 0.65 12 0.98 20 0.55 18 0.64 23 0.97 28 0.53 26 0.61 32 0.95 XII 36 0.51 34 0.59 37 0.95 41 0.50 39 0,58 40 0.94 46 0.60 43 0.58 54 0.94 69 0.48 57 0.56 65 0.93 70 0.47 67 0.56 78 0.92 84 0.46 81 0.53 96 0.92 100 0.46 97 0.62 117 0.92 121 0.44 119 0.50 144 0.91 < 3 M r i

t o k-* t o M *-* t-" t- J r o to ro r o k-k 4 H -J m A t o H* en t o (O g e n 4k r o i-* > -> ro t e 0 3 0 3 C g e n Cl* r o 1—< k~< t o 0 3 4 o 0 9 t o 0 3 "J t-> &4 - g 0 3 CD 0 3 g g 0 3 o r o 0 3 0 3 3-" r o 0 0 k-k en t-< 4k to 0 3 k4 g k-k g 0 3 a> en to Cl)

O o O O O O O O o O O o O O O O O o O O o O o o o O O O o O o o O o Q • 0 t o (O t o r o t o r o r o 4k 4k 4 k y n en en 0 3 0 3 (-1 t-3 M K* h-* k-* k-k M k-J M (-> ro 4k 4 4 4 4 4 1 t-" l-J M M t o CD en en g 00 o M 03 *-• CD en 0 3 03 03 0 3 0 3 03 03 0 3 g 00 ro 0 3 # 4 o * r o t o C4 k-k

w M M r o r o r o 1-3 1—' k J k-J Cl* Cl* t ) ro ror o k-k k-k k-3 < c '-3 9? Ok t o M O - 0 to O 00 g en 4k r o le) œ ch r o te 0 3 WO g en Ü»ro H* k-k •4 t o to 03 4 O œ 4 en O»o h-* en M o> m 1^ en 00 eo C4 (D r o 00 O

t-3 O O o O o O O O o O O o O O o O o CO O o O OO O OOO O O O O O O c o o OO O O Q • • O g* Oi 03 w o » 4 k en en en en en en 9 i 0 3 0 3 o> o>g g ro O) r o r o r o w Cl* 0 3 Cl) Cl* Cl* t •4 4 k 4 •4 4k 4 4 4 4 4 1 en en 00 r o 00 4k en 00 00 O* 03

X M t-j t-* to roto ro M (-> M Cl* C4 ro ro ro ro k-k g rt 00 'g en en O» ro !-■ o •g ■fk ro O 03 g en 4k ro «s g to (O en en to g ÎC W ro 3-* te Cl* to to 03 4 O g o 00 to o> 4k en en *• en en ro 03 g M g O 03 00 t£> ro V

9» »H r H to CM N •-* iH Q lA Q <0 o o» o » O) t o 0 0 g ) lO t a t a a O)OD OD O t o t o t o t o CM CM CM (D (O to <0 W O O 0> e n o t s < o «0 0 0 CD 0 0 g t - t - r H r H r H fH CM CM CD « 0 CD ODCD CD ***** * * «* • • • * * • •• • * # § r H r H r-t ( 1 r H r H O O O O O o o o o o o o o o o o o o o o — O o O OO O ■o

t - « r- CM to to to O CM CM O r H t o O » ■M* 0 > C M t o C- e n t o a a t o CD t " r H M* CD ^ 0> o Q C M ^ >-t CM (O «R t o t o C D o t o t o (H C" e n r H R r H v > OD CM t o CO O ^ o r H r H r H CM CM a r H r H (H 9 r H to f—f

t—i 0 CM to I—<05L0t0CMrHr-*Q C D C M t o t o t o CM CM CM O t o r H e n e n e n t o t o C M C M a e n rH 1 • o (O r H r H COtOtOtOtOtOtOtO t- c~ t o t o t o t o t o t o CM O O o 0 0 CD CD CD a a O ******** * O o O r H r H r H r H H oooooooo oooO O O r H r H oO O O oO

t o O) CMUïtotocncMaocM CD a> C M a e n r H CD CM « 0 t o ■M» t o a e n CMCM 6 0 t o t - CM l O r H t o t o CD O t o r H H f GO or H r H a 0 0 or H e-t r i CM r H r H r H r H CMCM r H r H r H r H

r l r H O kO CM en a a ts t e en a f a a CMCM r H CM en en en oa CM t -

CD r H OlOCM^mi— roOOQJllOCMtOO^CDCO co ai I—< CT> *0 tn lo LO en (O o CM t o ■HtOllît^OJCM^t^ CMIOirtt^CT>i-itOtOOOOtO r-t ^ to 0> o CM

Oi c ô o X (S» o g H G Table 13 (Contd.) ^7

Exp. No. T OD-1 T OD-2 T OD-3

17 1.06 16 1.07 14 1 .2 0 51 0.97 30 0.98 28 1.17 50 0 .8 6 48 0 .8 8 47 1.13 XX 66 0.81 66 0,82 63 1 .1 2 84 0.74 83 0.77 81 1.08 95 0.72 94 0.75 92 1.08 112 0.77 111 0.73 109 1.07 130 0 .8 8 150 0.69 128 1,06 147 0 .6 6 146 1,04 168 0.63 166 1.03 197 0.60 195 1 .0 2 219 0.67 217 1.02 302 0.98

14 0 .6 6 14 0.79 13 0.83 28 0.55 28 0,74 26 0.82 47 0.46 47 0 .6 8 46 0.80 63 0.36 63 0.62 62 0.76 xxr 81 0.35 81 0.62 80 0.78 91 0.31 91 0.60 90 0.77 110 0.50 109 0.58 108 0.76 128 0.33 127 0.56 126 0.76 144 0.54 143 0.75 167 0.52 167 0.74 194 0.49 195 0.73 215 0.48 214 0.72 307 0.54 300 0.71 372 0.69

15 1.13 17 1,21 8 1 .2 1 45 1.01 78 1 .00 56 1.18 XXII 61 0.95 120 0.95 58 1.15 80 0 .8 6 136 1.11 144 1.09 119 1.19 155 1.17 216 1.06 256 1.08 273 1 .1 2 291 1.15

15 0.76 15 0.80 17 0,87 49 0.57 78 0 .6 8 79 0.83 66 0.50 120 0.62 123 0.81 82 0.46 138 0.60 139 0.80 122 0.57 164 0.58 155 0.80 145 0,48 194 0.56 195 0.79 164 0.49 216 0.64 218 0.78 236 0.723 237 0.78 108

T&bl# 13 (Contd*)

Exp. No. T OD-1 T OD-2 T OD—3

16 1.19 16 1.22 17 1.22 m v 37 1.14 58 1.17 59 1 .2 1 59 1.08 61 1.14 62 1.19 80 1.56 81 1.29 83 1.18 110 1.40 112 1.17 144 1.14 IT# 1.19 206 1.28

21 0,85 25 0.83 25 0.84 44 0.81 46 0.85 48 0.84 65 0.79 67 0.82 69 0.84 86 0.76 89 0.81 91 0.83 XXV 116 0.74 117 0.80 119 0.83 148 0.96 151 0.80 166 0.82 185 0.80 187 0.82 208 0.79 211 0.82 247 0.78 249 0.83 277 0.78 280 0.82 292 0.78 293 0.82 309 0.78 315 0.79

11 1.18 15 1 .2 0 16 1,22 46 1.09 48 1.14 51 1.19 77 1.02 79 1.08 81 1.16 107 0.962 109 1.04 111 1.15 153 0,91 156 0.99 137 1.12 156 0,87 157 0.96 160 1.11 XXV7 179 0.82 181 0.83 183 1.08 200 0.80 201 0.91 204 1.08 216 0,80 218 0.89 221 1.07 220 0.80 261 0.94 263 1.0 ? 223 0.81 285 1.02 319 0.85 297 1 .0 1

32 0.75 49 0.79 20 0.84 49 0.70 67 0.76 50 0.81 68 0.65 92 0.74 79 0,80 93 0.60 116 0.72 97 0.79 xrvii 116 0.56 160 0.69 140 0.77 140 0 . 6] 176 0 .6 6 163 0.76 160 0.48 279 0.60 187 0.75 176 0.45 301 0.59 207 0.75 109

Table 15 (Contd.)

Exp. No. T OD-1 T OD-2 T OD—3 r m i 220 0.40 513 0.58 224 0.74 (Contd.) 241 0.46 527 0.71 349 00.70 361 0.70

15 1.20 15 1.22 17 1.22 55 1.16 37 1.19 39 1.21 67 1.14 69 1.17 71 1.19 91 1.11 92 1.14 94 1.18 XXVIII 108 1.08 109 1.12 111 1.16 146 1.04 150 1.08 152 1.16 176 1.01 177 1.07 160 1.15 216 0.96 218 1.05 220 1.15 248 0.95 250 1.00 252 1 .1 2 269 1.06 271 0.99 273 1 .12 299 1.12 301 1.11 316 1.09 357 1.09 369 1.08 591 1.08

19 0.82 21 0.85 24 0.84 42 0.79 44 0.82 46 0.83 75 0.74 77 0.80 79 0.82 97 0.71 99 0.78 100 0.81 XXIX 115 0.69 115 0.78 118 0.81 155 0.64 156 0.77 158 0.80 182 0.61 184 0.75 186 0.80 223 0.57 224 0.74 226 0.80 254 0.57 256 0.72 268 0.79 275 0.67 277 0.71 279 0.79 506 0.70 307 0.78 320 0.69 322 0.78

19 1.21 23 1.19 16 1.21 39 1.20 42 1.19 37 1 .2 0 60 1.18 64 1.17 69 1.19 XXX 114 1.15 lie 1.14 91 1.18 135 1.15 156 1.14 107 1.17 198 1.19 202 1.09 126 1.16 147 1.16 202 1.14 222 1.14 286 1.11 110

T able 15 (C ontd.)

Exp. No. T OD-1 T OD-2 T OD—3

27 0.85 19 0.83 49 0.84 47 0.82 40 0*82 97 0.82 67 0.82 73 0.82 111 0.82 122 0.80 95 0,81 151 0.82 144 0.79 lie 0.81 151 0.82 211 0.78 129 0.80 206 0.81 150 0.80 228 0.81 204 0.80 295 0.81 227 0.80 295 0.80 CM to h C to CO CMCMCM A œ aac- to to to o> 0) a* o> A AA to to to tos to $tO to é ooO oO OOO oO ooooooo

fi o U} <£) f-( lO ^ lO dO lOC'IOO^ODl^OC'OJ -p n lO CO CO o 00 rHtOiOC'CBrH'M'C'rH P* rH rH f—4 CM to « A AAA a> #o ■QCMCMrH rH A * •0 tc w lO •O•0 « ■C CM -p sss c o o o oO ooooooooo oo •

1 COaOCMCOOD'M'Ot^ tOlOtOrHC-t-^Qt'lO s rM to CO Q »0 ® C' CMHfC-cOCMtOdDOCM ■ A o o A to to to A n A A A A rH A œA A ■P -* ? M* ^ -Mt A A AA A 00 A A m A A A C- •H • • • B oo O C oo o oooO o O oooOO ooO Ô o • A rH B tOO^C-rHiOeO^CM ^nCtrHrHt-iCtOQtOeCrHcnt^ O Sj^b-rHiOtOeOcS rHIOtOC^O»rHM)OOOIOtOcOO) *#H rH rH rH rH PS i-HrHi—IWCMCMCMIO I lai 112

Tkbl* 16

Optioal Density of KllnO^ in Vertovs Salt Solution*

Exp. Exp. No. T OD-1 T OD-2 No. T OD-1 T OD-2

9 0.82 9 0.84 14 0.76 15 0.78 26 0.79 26 0.85 27 0.74 26 0.76 45 0.76 44 0.82 45 0.72 44 0.76 X3CXV 61 0.73 60 0.81 XXXVI 61 0.69 60 0.74 77 0.71 76 0.80 78 0.67 77 0.74 94 0 .6 8 93 0.80 94 0 .6 6 93 0.74 109 0 .6 6 108 0.80 114 0.64 113 0.75 158 1.07 158 0.78 164 0.60 165 0.75 202 0.77 208 0.93 207 0.71 261 0.76 267 0.71 505 0.75 509 0.70

14 0.82 14 0.85 15 0.71 14 0.72 30 0.79 50 0.83 50 0,69 29 0.71 52 0.76 62 0.82 62 0 .6 6 51 0,98 XXXVII 78 0.71 78 0.81 XXXVIII 78 0.63 77 0 .6 8 101 0.76 100 0.80 102 0.61 101 0 .6 8 140 0.78 140 0.58 139 0.67 168 0.77 168 0 .6 6 167 0.67 119 T able 16

Optical Tensity of KMnO^ in Various Salt Solutions

T OD-1 T OD-2 T OD—5 T OD-4 T OD—5 T OD-6

6 0 ,6 6 5 1.09 7 0.64 6 1.05 6 0.79 6 1.09 17 0.61 16 1.00 17 0.60 16 1.03 16 0.72 16 1.07 27 0.60 26 0.99 27 0.67 26 1.02 51 0 .6 8 32 1.06 38 0.57 37 0.99 42 0.56 41 1.02 46 0.67 48 1.04 64 0.56 53 0.97 58 0.54 67 1 .00 63 0 .6 6 64 1.04 TO 0.55 69 0.96 74 0.55 73 0.99 89 0.64 90 1.04 xxxrx 85 0.55 84 0.95 100 0.52 99 0.98 111 0.65 115 1.03 111 0.62 110 0.94 123 0.51 122 0.97 155 0.61 156 1 .0 2 154 0.52 135 0.94 166 0.50 165 0.97 187 0.60 188 1 .0 2 177 0.50 176 0.98 198 0.50 197 0.97 225 0.59 226 1 .0 1 209 0.50 208 0.95 236 0.49 236 0.96 264 0.58 266 1 .0 1 247 0.49 246 0.92 276 0.49 274 0.96 304 0.56 506 1 .00 286 0.47 285 0.91 315 0.48 514 0.94 326 0.46 325 0.92

21 1.15 20 1.15 19 1.16 14 1.18 13 1.18 12 1.17 35 1.10 34 1.09 53 1.11 35 1.10 54 1.09 35 1.07 54 1.05 53 1.03 52 1.04 53 1.04 52 1.03 51 0.98 69 0.99 68 0.98 67 1 .0 0 68 1.02 67 0.97 66 0.95 86 0.94 86 0.95 64 0.95 TO 1,00 69 0.97 68 0.92 105 0.89 104 0 .8 8 103 0.90 86 0.96 85 0.92 84 0.87 121 0.65 120 0.84 119 0 .8 6 104 0.91 103 0.87 102 0.81 157 0.81 156 0.80 125 0.82 120 0.87 119 0.82 118 0.77 136 0.84 156 0.78 134 0.72 277 0.63 276 0.48 275 0.45 114 1.30

9 MNO, Run O HCIO4 Run D ^^SQ, Run 120 A kyO , Run • HF Run

I.IO —' {M n0;)-50 « KT^Mo* run* 1.25 K K)"’ Mall run* Acid Cone -3.1 N oil runt

100

0.90

>K

0.70

0.60

0 5 0

0 40

0 30 O 30 60 90 120 150 100 210 240 270 300 Tim# (Minutât)

Figure 20. Comporison curves for the system Mn*'*', MnO^ , acid 115 T 1------1------T

0.60

0.70

8 a. o I &5mlO-T6 KMnO* O l.75itO'^M MnSOu

r llM HOQ* 5A» K)-*M KMnCU

0.60 12.0* HCKk A 5Axl0-> KMnO# I l.75mcr*M Mnt^

2.0 M HCK& &5« 10 M KMnO* [3u5>KT*M MnSOW [5.IM CtCf,2.0M MCIQ, O &5*Kr% KWMCW Ll.75mKr*M MnSC^

5.l M C*Q7 ^.OM HCIO# 0.50 SwSaKT^ KMnO# rL&6 %X)-*M NMtO* I I I I 120 160 200 240 260 TWn# (Minutât) Figure 21. Soit effect* on ttie decomposition of KMnQ^ l0 9 (MnGC) -4.000 -3.000 -2.30

-2.40

-2.50 V (Mn**).S.OOmKr'M -2.60 Û tt(in<Ç)-2.00i KT^M □ (NMî)-9l50m KT^M -2.70 O (&W%)"5jOOm KT^M (HCK^)- 3.1 M for a l M -2.80

-3.00

-3.10 -

-3.40 A ^

-3.60

-3.70 •---- -4.500 -4000 -3.000 -2.500 log(Wo**)

Figure 2 2 . Variation of the order of reaction with coocentrotion. SUittAKY

The problem in this rese&roh Ttaa to detersiine whether an iiotcpie exchange ooourred between manganoua ions and permanganate ions in acid solution and to determine some of the parameters of this exchange process. Radioaotire manganous ions were used as tracers. Even though the results obtained by previous experimenters indicated exchange between these two ions, their results were uncertain because of the rapid appearance of intemediate oxidation states of manganese in this system.

The concentration of these intermediates was found to be in the order of 10 to 15 per cent of the total concentration of manganous ion a few minutes after mixing as Indicated by spectrophotometric data.

Therefore, in order to establish the validity of the data of previous experimenters and of this research it was necessary to determine whether these intermediates were separated in the manganous ion fraction or the permanganate ion fraction. If these intermediates appeared in the permanganate fraction, then the activity observed in this fraction could have been due to activity in these intermediates rather than in the permanganate ions. If this were true, no exchange would have occurred between Mn*^ and àlnO^.

The previous investigators were aware of the presence of these intermediates but apparently made no attempt to consider the effect of their presence upon the cata obtained. It was the concern of this research to develop exerpiments which would take into consideration and account for the wh «raboute of these intermediates and to establish

1 1 ? 118 the vmlidlty of the exohang;e. Thia eea accompliahed in part by using two different procedures for separating the MnO^ tagged ions

formed from the initially tagged Un ions.

The intermediate oxidation states of msinganeae expected in the

system Un**-MnO^-aold and some inform ation about th e ir s t a b i l i t y is as

follow s;

The hypomanganate ion (MnO^ ) is very unstable in acid solution 4+ « a and di sproportionates quiolrly, presumably into Un and UnO^ • MbO^ is

also very unstable in acid solution and presumably dl s proportions te into — 3* “ UnC^ and UnO^ * ihe UnO^ ions are much more stable in acid solution than the other ions and therefore serve as a stabilising species In the a 3» disproportionation of MnO^ and MnO^ ions. These latter ions may be

considered present in very low conoentretions only. The same is true 4 + for the Mn since any appreciable concentration of these ions in

solution results in the rapid formation of manganese dioxide. Thus it

appears that the only intermediate oxidation state of manganese that

can exist in relatively high concentration in acid solution is the +5

s ta t e .

+ + — It was found in the system Mn -UnO^-EF that the predominant intermediate oxidation state of manganese was the Mn species which

supports the statement made in the previous paragraph. The determination was made using an ion exchange re sin separation method. I t was assumed that the +3 state was the main intermediate oxidation state in. other acid systems (nitric, sull'urio, phosphoric and perchloric) stuoied in this research. 119

Since this iatermediete rtate (Ifn^^) does exist in appreciable

concentration in Mn**-MnO“ solutions, it was necessary to detem lne 4 what happens to this ion in this +5 state during the separation procedure

in order to answer the question, "Does the Mn*® species precipitate with +♦ * the Mn ions or do they remain in solution with the UnO^ ions and thus

account for the aotlvity observed In this fraction?" To solve this

problem the two separation procedures mentioned previously were carried

out along with other supplementary experiments.

In the first procedure, sodium hydroxide vas added to the system

Ma**-Mn 04“HC10^ until the solution was slightly basic. In this basic

solution the Mn** ions wore rapidly oxidised by MnO^ ions to manganese

(IV> which then p rn c ip ita ted as manganese dioxide. The p re c ip ita te was

separated from the solution containing the permanganate ions. Each

fraction was then analysed for manganese using the silver oxide method.

In working out this method of analysis and senarstion it was

observed that the longer the Mn ions remained with the MnO^ ions in

acid solution the greater the amount of Mn** oxidised to MnOg(which precipitateo^ This may pcssibly be accounted for as followsi

An increase in concentrât ; on can occur in this fraction (WnOg) by preoipitntin; the intermediate oxidsti on states which arise from

■+ + w oxldet: on ana reduction of both the Mn and mnO^ ions. This can be 4 seen by considering the following reactions which may occur on n e u tra lis a tio n of the Mn^^-MnO^-ac:d so lu tio n si 120

(l) + 2 M1 1 O4 + 2 H2 O = 5 MnOg + 4 H (Guyard Reaction)

(II) 3Mn"*"*^ + MnOj + 4 HgO = 4 MnOg + 8 s'*

++ (III) Mn + 2 OH • Mn(OH)g

(IV) Mn**"^ + 3CE" * MnO(OH) + HgO lieta from th e experiment» indicated th a t the Guyard reac tio n waa taking place in the soluticna. It is seen from the above reactions that the

Guyard reaction produces the highest concentration of manganese in the precipitate. If this reaction is taking place in the solutions then, in order to get a higher concentration of manganese in the precipitate, it is necessary to obtain manganese which was originally in the form of permanganate ions, that is, the intermediates formed from the decompos­ ition of the permanganate ions. The increase in concentration of man­ ganese actually observed in the precipitate corresponded closely to the decrease in concentration of permanganate observed spectrophotometrically.

Thus it appears that the intermediates do precipitate with the manganous ion fraction.

To supplement the above results, an isotopic exchange experiment was carried out in hycrofluoric acid solution. No isotopic exchange was obsf.rved in this system (Mn**-MnO^-HF) . It was observed in this research the main intermediate oxidation state in the system was the

+3 s ta te . I t was found, upon addition of sodium hydroxide to the system, that a precipitate formed idiich centalned over 98 per cent of the activity. Therefore it appears that the +3 state was almost com­ pletely removed. If any remained in the solution, its concentration 121 would not be nearly high enough to account fo r th e a c tiv ity observed in the MnO^ fraction.

The most favorable conditions under which the +3 state would be lik ely to remain in the MnO^ fra c tio n would seem to be i f th a t state were strongly oomplexed. Since the +3 state was strongly complexed in the hydrofluoric acid solution (compared to the other aoids used) and yet the +3 state was almost completely removed, it would appear th a t the s ta te would be removed in other less strongly complexing acid solutions also.

In the second procedure, Dowex 1-X8 anion exchange realn was used to separate the MnO^ ions from the Mn*^ ions. It was observed in previous experiments that the Mn ^ and Mn ions passed through the resin quite readily, however, permanganate irns were strongly absorbed at the top of the column. Therefore this resin served as a means of separating the lower oxidation states of manganese from the +7 state.

In passing a solution containing Mn*^, MnO^ ions (and intermediates) and perchloric acia througii the resin it */as found that no separation induced exchange occurred which indicates that there was very little reaction with the resin ana therefore a reletively clean-cut separation of the lower oxidation states from the permanganate ions. If this is the case then any activity observed in the nermanganate fraction must be due to activity in the MnO^ ions and not to activity in the inter­ mediate oxidation states since tliose passed through the resin.

Activity was foi.nd in the permanganate fraction from the resin. In fact the isotopic exchange results obtained from these resin experiments 122 agreed with those obtained using the sodium hydroxide method discussed above. Thus it appears that the Intermediates are not responsible for the activity in the permanganate fraction.

from these ex[jieriments it may be concluded that a major part of the intermediate oxidation states are removed with the Kn** Ions in the separation procedure. Therefore, it follows that a true exchange

+ + — does coour between Mn and MnO^. Concentration cf the intermediate oxidation states remaining in solution are too low to account for the higli a c tiv ity observed in th e MnO^ fra c tio n .

The last section cf this research deals witi. the reduction of permanganate in the presence of manganous ions. The rate of reduction was followed spectrophotometricelly.

The spectrophotometri c studies si owed a continuous decrease in permanganate concentration in the presence of manganous Ions. Thus, a continuous increase in concentration of intermediate states occurred in the system. The time of appearance of manganese dioxide def>ended on the initial concentration of Mn*^ and MnO ions; higher concentretions 4 resulted in more rapic. appearance of the p re c ip ita te .

The spectrophotometri c studies were mace in order to determine some of the parameters tliat effect the kinetics cf the system. It was determined from the data tnat the order of the reaction with respect +♦ to Mn changed with changing initial Mn ion concentrations over a certain concentration range while the order of the reaction with respect to MnO^ remained essentially constant for all permanganate ion concentrations. From the rate eouation 125

-d c/d t “ k(Mn**)*(MnO^)^ i t was founü. that ttie value of ”a" was about 1.6 in the higher concen­ trated solutions in which the manganous ions were in excess, while in the solutions in which the permanganate was in excess the «^lue of "a" was approximately 0.67, The value of was about 0.85 in all solu­ tions. A more complete discussion of the kinetics was given in the previous section.

Due to the changing order of reaction and complexity of the system no d e fin ite mochanism can be given th a t w ill cover a ll conoent ra tio n ranges. However, from the study of the complexing ability of the manganous ions (which essentially eliminated the +2 state in the rete- determining step) it appears that, the rate-determining step is between the +3 ana *A states as indicated in Adamson's work. This step apparent­ ly involves a change in configuration from a six coordi nateo structure / 3* » 2" (e*g* MnXg ) to A lower coordination nuinbf r, poesibly MnOj • This conversion could account for tliis step being slow and therefore, the rate-determining step.

Yore work has been done on the k in e tic s of th is system in the previous section.

Conclusions

+ + — (IJ An exchange does occur between Mn and MnO^ in acid solution by means of an electron transfer mechanism. (2) The half- time for the exchange is approximately 160 minutes in an atmosphere 124

of air or nitrogen; in oxygen the half-time is approximately 100 ndnutee.

( 3 ) Nearly all of the intermediate oxidation states precipitate with the Mn** fraction* (4) Msoigancus ions form weak complexes with the * ** * • +■ + anions CIC^, NOj, SO^, F and and therefore the Mn ions are

assumed not to enter into the rate-determining step. The rate-determining

step appeared to be between Mn^* and Mn^*. (5) Kinetic data show that

different reactions predomi nate in Mn**-i.'nC^ solutions according to the

ion in excess and the concentration range covered. (6) Perchlorate and

nitrate ions have no noticeable effect on the decomposition rate of

potassium permanganate in acid solutions. The sulfate ion has some effect

on the rate. The di-hydrogen phosphate ana fluoride ions have the

greatest effect on the rate. AUTOBIOGRirar

I , B a ri Lawla Iferrynan* KM b o m In C&ditr« O hio,

Jan u ary 3 0, 1931# I reoelTod my #eoomdary—aohool education in the public achool of St« C lalraville, Ohio, and my under* graduate training at Ohio State Unlvereity, mhlch granted me the Bachelor of Science degree in 1956. I remained at Ohio

State University for my graduate irork. While in residence there, I was an assistant in the Physical Chemistry laboratory during the academic years 1956*57 and 1957*58, and assistant* instructor during the academic year 1956*59# I #&s granted the

American Cyanamld Fellowship in October cf 1959 and a

Rational Science Foundation Fellowship from June to August of i 9 6 0 . I am at present employed at Battelie Memorial Institute in Columbus, Ohio, and at the same time completing the requirements for the Doctor of Philosophy degree*

125