ABSTRACT

The rate constant for the Tl(III)-re(phen)§+ reaction
has been measured for thalliua(III) concentrations greater than

“.7“ x 10.3 g at”: 3.0. The value for the rate constant at 25°C.

2 liter sole-lsec'lg at 30°C. the rate constant

1

1. “079 1 10.
is approximately 0.528 liter mole- sec‘l. The reaction appears

to be second order.

d[?e(phen)§f7
dt

 

. x [Fe(phen)fl('1‘13+)

At thalliun(III) concentrations less than h.7# x 10"3 g
the iron phenanthroline complex dissociates rapidly, and the
oxidation reaction occurs simultaneously. When lanthanul(IlI)

was substituted for thallius(III), it was noted that

J.

 

d F ( h )2”
‘EO P en 1. 1‘1 4. kzfi‘dphengfl;

dt

.2

The rate constants were calculated: k1 is 8.87 x lO sole.

liter' sec' and k2 is 5.95 x 10 sec .

A KINETIC STUDY OF THE
THALLIUM<III)-TRIs-(1,1o-PHENANTHR0LINE)IR0N(II) ION

RCACTION IN NITRIC ACID SOLUTIONS

by

Penny Jean Stoner

A THESIS

Submitted to Michigan State University for the degree of
Master of Science

Department of Chemistry
1967

DEDICATION

To Mom and Dad

ACKNOWLEDGEMENT

I wish to thank Dr. Carl R. Brubaker, Jr. for the
many helpful sugaestions and the encouraqement he gave

during the completion of this project.

TABLE OF CONTENTS

page
Abstract h
Introduction 5
Historical 6
Preparation and Standardization of Stock Solutions 10
Experimental Procedure 16
Results and Discussion 18
Appendix 1 35
Appendix 2 36
Appendix 3 37

Appendix 4 38

LIST OF FIGURES

title
1. [Fe(phen)2;J
Time dependence of In ——-3+—L in perchlorate
. (T1 )
solution..
2 . G‘s (phen )2?
Time dependence of In ——3:-2- in nitrate
(r1 )

solution.

11

page

20

22

1.

2.

3.

The
for

ion

The
for

ion

LIST OF TABLES

effect of {Fe(phen)§£]on the rate constant
the Tl(III)-tris-(l,lO-phenanthroline)iron(II)
reaction at 25°C.

effect of[:Fe(phen)§flon the rate constant
the Tl(III)-tris-(l,lO-phenanthroline)iron(II)
reaction at 30°C.

Tris-(l,lO-phenanthroline)iron(II) ion- A1(III),
Ce(III) or La(III) initial-rates.

Tris-(l,lo-phenanthroline)iron(II) ion-Sr(II)
initial rates.

Rate constants for the La(III)-tris-(l,lO-

phenanthroline)iron(II) ion reaction.

Hydrogen ion dependence of the Tl(III) and La(III)
reaction with tris-(l.lO-phenanthroline)iron(II) ion.

111

Disc

21

23

24

25

26

29

INTRODUCTION

A study of the Thallium(III)-Tris-(l,lO-phenanthroline)
iron(II) system at high thallius(III) concentration indicated
that an oxidation reaction which was first order in each of
the reactants occurred in perchlorate solutions. It appeared
that a similar result should be obtained using lower thallium
concentrations.

It the nitrate radiacal did not cause a specific anion
effect, a result analogous to that obtained for the perchlorates
should be discovered in nitrate solutions. A series of experi-
ments with varied thallium(III), iron-phenanthroline complex
and hydrogen ion concentrations was performed to determine the

course of the reaction and its rate in nitrate solutions.

HISTORICAL

The first studies on the iron-phenanthroline system
performed in 1936 showed that the intense absorption bands of
the complex in the visible region were readily distinguishable
from the ligand bands appearing at approximately 270 my“ which
were not due to the electronic transitions of the specific
cation.(l)

Kolthoff, Lee and Leussing(2) conducted equilibrium
studies on tris-(l,10bphenanthroline)iron(II) complexes and
determined that the equilibrium constant, K', was 5 x 10'22

for the reaction,

2

Fe(phen)§+ I Fe + + 3 phen.

For the dissociation in acid solution the equilibrium constant,

K", was A x 10-7 for the reaction,

2+

Fe(phen)§+ + 3 If u Fe + 3 phenHI.

The calculated rate constant for the dissociation reac-
tion in 1 1; acid was 7.5 x 10'5 aeo'1.(2,3) The half-life in
acid solution was approximately 2.5 hours at 25°C.

In acid solution phenanthroline adds only one proton
due to steric factors. The mono- and the bis- as well as the
tris-(l,lO-phenanthroline)iron(II) ions are present in appreciable
concentrations. In neutral solutions only the tris-(l,lO-

phenanthroline)iron(II) concentration is appreciable.(#)

7.
Basolo, Hayes and Neumanm(3) dbtermined the dissociation
constants for tris-(l,lO-phenanthroline)iron(II) ion in nitric

acid. Their results are as follows:

Temperature (°C) ‘ - kdiss (800-1)
25.3 7.38 x 10-5
30.9 1.97 x 10'“
34.6 3.70 x 10‘“

Increasing the concentration of nitric acid from 0.01 g

-1

to 1.79 11. causes a decrease in kdiss from 7.57 x 10'”5 sec to

6.13 x 10"5 sec-1

. When potassium nitrate is added, the rate
constant increases with increasing potassium nitrate concentra-
tion. The addition of increasing amounts of lithium nitrate
caused a decrease in the rate constants.(5)

The first kinetic studies of the Tl(I)-Tl(III) exchange
reaction were conducted by Prestwood and Wahl(6) by using
radioactive Tl(III). Perchloric acid solutions with sodium
nitrate added to regulate the ionic strength were used. The
rate of the reaction varied with the acid concentration and with
the acid used. The rate was greater in nitric and hydrochloric
acids than in perchloric acid. The maximum rate in each acid
was observed at l 5 acid concentration. The addition of lithium
perchlorate to perchlorate solutions increased the rate of
exchange.

The Tl(I)-Tl(III) system was first order in total con-

centration of thallium(I) and first order in thallium(III).

8.
The rate of exchange decreased with an increase in hydrogen ion
concentration. Both Tl(III) and TlOEaI were considered present
in significant concentrations, but only T1082+ was assumed to
exchange with Tl(I).(7) ‘

In the Fe(II)-T1(III) system the rate constant decreased
when the hydrogen ion concentration was increased from 0.75‘fl
to 1.5 g, When the hydrogen ion concentration was increased from
2 .13 to it E, the rate constant increased. The rate constant de-
creased slightly after the reaction was 50-70% complete. In the
absence of oxygen this decrease was not noted.(8)

Johnson(9) studied the same system. Hydrogen ion concen-
trations large enough to be considered constant throughout each
run were used. The rate constant was identical for excess iron
and excess thallium. Hydrolysis products of thallium were assumed
to be important in the rate determining step.

The pH dependence of the electron exchange between Tl(I)
and Tl(III) in perchloric acid solutions was studied by Born,
Vogg and Vogt(lO), who postulated that an increase in hydrogen
ion yielded complex formation between Tl(III) and perchlorate
radicals. The rate constant equalled 3.33 x 10"5 sec'l.
I Duls and Sutin investigated the kinetics of the tria-
(l,lO-phenanthroline)iron(II)-Ce(IV) reaction. The reaction was
first order with respect to each of the reactants. The rate
constant depended directly on the hydrogen ion concentration,

but it was independent of the initial concentrations of cerium

(IV) sulfate and tris-(l,lO-phenanthroline)iron(II) sulfate

9.
over a wide range of concentrations.(ll) The acidity effect was
samll. Sulfuric acid from 0.5‘fllto 0.05 fl_was used.(12)

The rate of reaction between T1(III) and tris-(l,10-
phenanthroline)iron(II) ion in 2.7 fl'perchcoric acid was ob-
served by George and Irvine(l3) in l95#. The rate constant for

the reaction was approximately 7.3 x 10.3 liter mole"1 sec-1

at 25°C. when the T1(III) concentration was 3.6 x 10-2 fl and
the tris-(l,10-phenanthroline)iron(II) ion concentration "as

5.0 x 10'5 g.

PREPARATION AND STANDARDIZATION OF STOCK SOLUTIONS

1. Nitric acid.

Concentrated nitric acid was dissolved in demineral-
ised distilled water. The solution was standardized by titra-
tion with standard, carbon-dioxide-free sodium hydroxide.
Methyl red was used as the indicator.

Perchloric acid solutions were prepared in the same

manner.

2. Tris-(1,10-phenanthroline)iron(II) nitrate.

Iron(II) nitrate nonahydrate(Baker Analysed) was recrys-
tallised three times from demineralized distilled water. Iron
(II) nitrate and three times as many moles of l,lO-phenanthro-
line were dissolved in dilute nitric acid. The trio-(1,10-
phenanthroline)iron(II) nitrate was precipitated by the addi-
tion of concentrated sodium nitrate and cooling in a freeser.
The solution was filtered, and the precipitate was dissolved in
demineralised distilled water. The perchlorate salt of the com-
plex was prepared in a similar manner.(14)

The concentration of the complex was determined spec-
trophotometrically by use of a Beckman DU Spectrophotometer
and published values for the molar absorbtivity. The wave-
length of maximum absorbance is 510 If. and the molar absorb-

tivity at this wavelength is 11.1 x 103.(15)

11.
3. Tris-(l,10-phenanthroline)iron(III) nitrate.

(a.) Tris-(1,10-phenanthroline)iron(II) nitrate was
oxidised by the addition of an excess of cerium(IV) sulfate.
(5) '

(b.) Tris-(l,lO-phenanthroline)iron(II) nitrate was
dissolved in dilute sulfuric acid and oxidised with an excess
of lead dioxide. The excess lead dioxide and lead sulfate were
filtered off. The filtrate was placed in a freezer to precipi-
tate the iron(III) complex. The complex was filtered out of the
solution ahdflthén dissolved in a small amount of 1'g nitric acid
and reprecipitated in a freeser. The precipitate was dissolved in
9‘fl’nitric acid and stored in a freeser in a brown bottle.(lh)

Perchlorate solutions of tris-(l,lO-phenanthroline)-
iron(III) ion were prepared in an analogous manner.

The complex concmmtration was determined spectro-
photometrically at 600 mp. The cerium oxidised solution cannot
be standardised at that wavelength because cerium(IV) also
absorbs at that wavelength. The molar absorbtivity of the com-
plex at 510 meis 300.(5) There is no interference from cerium

(IV) at the latter wavelength.

A. Thallium(I) nitrate.

Thallium(I) nitrate(Fisher) was dissolved in hot demin.
eralised distilled water, and it was recrystallized when the
solution was cooled. The process was repeated twice. The pre-

cipitate was dissolved in dilute nitric acid.

Standardization was accomplished by adding hot concen-
trated hydrochloric acid to a measured sample of the T1(I)
solution to make the sample 5-8% hydrochloric acid. The acid
was heated(no-7o°c.) to prevent the precipitation of the
thallium(I) chloride and to prevent clogging of the saturated
calomel electrode tip. The sample was then titrated poten-
tiometrically with standard potassium bromate. A Beckman

Model G pH meter with saturated calomel and platinum electrodes

was employed.(l6)

5. Thallium(I) perchlorate.

Thallium(I) nitrate(Fisher) was dissolved in hot con-
castrated perchloric acid and heated at 105-13500. to remove
the nitric acid formed. Thallium(I) perchlorate precipitated
when the solution was cooled. The precipitate was recrystallized
twice from dilute perchloric acid.(17) The filtrate gave a neg-
ative brown ring test for nitrate. The precipitate was dried
at 120°C. for an hour. It was then weighed and dissolved in
dilute perchloric acid. The solution was standardized by the

procedure described for the thallium(I) nitrate.

6. Thallium(III) nitrate.

Thallium(I) nitrate(Fisher) was dissolved in concen-
trated hydrochloric acid. Thallium(I) chloride precipitated
and was suspended in 1.! hydrochloric acid.and heated to 60°C.

An equivalent amount of potassium bromate was added to the

13.
hot solution to oxidize the thallium(I) to thallium(III). The
solution was made basic with concentrated ammonium hydroxide.
The thallium(III) hydroxide precipitated and was washed with .
dilute ammonium hydroxide by decantation until there was no
evidence of chloride ion when a sample of the wash solution
was added to silver nitrate and nitric acid.(18) The precip-
itate was dissolved in 9 fl nitric acid and stored in a brown
bottle in the dark. Thallium(III) perchlorate was prepared in
an analogous manner.

To standardize the solution a measured sample was re-
duced to thallium(I) by bubbling sulfur dioxide through it.
The sulfur dioxide reduced the nitrate first. When all the
brown vapors were expelled, sulfur dioxide was bubbled into the
solution at a moderate rate for an additional twenty minutes.
The solution was then boiled for an hour to remove any excess
sulfur dioxide(22) and titrated according to the procedure for
thallium(I) nitrate.

To determine the hydrogen ion concentration in the
thallium(III) nitrate stock solution an one-hundredfold excess
of sodium bromide was added to complex the thallium before the
solution was titrated with standard sodium hydroxide by using

methyl red indicator.(l9)

7. Sodium nitrate.

Sodium nitrate(Baker Analyzed) was recrystallized twice

and then dissolved in demineralized distilled water. After

if:
being filtered through a fine fritted funnel, the filtrate was
boiled gently until small particles appeared on the surface of
the liquid. A small amount of water was added to dissolve the
solid before the solution was placed in a 60°C. oven for twenty-
four hours. The crystals were filtered and left unwashed.(l6)
Sodium perchlorate prepared in this manner was obtained from
James Hoeschele.

The sodium nitrate crystals were dissolved in deminer-
alised distilled water, and the solution was standardized by
placing measured aliquots in weighed porcelain crucibles with
lids and drying for twenty-four hours at 165°C. with the cru-

cible lids askew. The crucibles were then covered and cooled

for thirty minutes in a dessicator before weighing.(l9)

8. Potassium bromate.

Potassium bromate(Baker Analyzed) was dried in an oven
for twelve hours at 165°C. and then were allowed to stand over
concentrated sulfuric acid for an hour before weighing. The
weighed potassium bromate was dissolved in demineralised dis-

tilled water.

9. Lanthanum(III) nitrate.
C.P. Lanthanum(III) nitrate(Fairmcunt) was dissolved in

demineralised distilled water.

10. A1uminum<III) nitrate.

Reagent grade aluminum(III) nitrate(Baker Analysed)

W

was dissolved in demineralised distilled water.

15.
ll. Strontium(II) nitrate.
Reagent grade strontium(II) nitrate(Baker Analyzed)

was dissolved in demineralised distilled water.

12. Cerium(III) nitrate.
Cerium(III) nitrate(Fisher), reagent grade, was dissolved

in demineralised distilled water.

EXPERIMENTAL PROCEDURE

A thermostated Beckman DU Spectrophotometer was employed
to obtain all spectrophotometric data. The kinetic studies were
conducted in the presence of oxygen as initial studies in a
nitrogen atmosphere showed no variation in the results.

All the reagents used in the kinetic experiments were
stored at 25°C. in a thermostated bath for one hour prior to
each experiment. The moment at which the tris-(1,10-phenanthro-
line)iron(II) nitrate or perchlorate was injected into the
sample cell was recorded as the intial time, to.

The experiments were conducted by utilizing trim-(1,10-
phenanthroline)iron(II) concentrations which were varied within

1’ g to 6.19 x 10"6 g. The thallium(III)

2

the range 1.21 x 10-

concentrations were varied between 3.90 x 10-

10-6 n...The hydrogen concentration range was 2.9 g to 0.5 a.

£1 and 2.11 x

In preparing the kinetic solutions the thallium(III)
was introduced into a four milliliter spectrophotometer cell,
one centimeter in width, followed by the addition of the acid.
The iron-phenanthroline complex was added at the moment the
timing of the reaction was begun. This moment was thirty seconds
after the addition of the acid.

To eliminate the possibility of any disappearance of the
complex before the thallium(III) was added to the kinetic'i ,
solhtion,,the'aciliwas added to.thhsthsllium(IIIITsolition‘

which was intially 0.5 g in hydrogen ion.

12,.
The final volume in the sample cell for each experi-
ment was three milliliters.
I The absorbance of the tris-(l,10-phenanthroline)iron(II)
ion at 510 mp.was recorded at various time intervals until
approximately 90% of the original iron(II) complex had dis-

appeared.

RESULTS AND DISCUSSION

An experiment at the concentrationsutilized by George
and Irvine, thallium(III) perchlorate s 3.6 x 10.2 E, tris-
(l,lO-phenanthroline)iron(II) ion a 5.0 x 10"5 g and hydrogen
ion a 2.7 5, had a rate constant equal to 7.32 x 10.3 liter
mole'l sec-1. The rate constant reported by George and Irvine
(1h) was 7.3 x 10"3 liter mole'1 sec-1. The reaction appears to
be first order in thallium(III) and first order in tris-(1,10-

phenanthroline)iron(II) ion.(Fig.1)

 

2
- dfi‘dphenbfl s kE‘e(phen)§q (1313*)

 

dt
giy.(phen)2:1 {Te(phen)2‘] - 2(T13I) h ‘Te(phen)2;1o
1._________2__£ . 4—2 ° ° kt + ln'----;z--
(r13.*)t 2 (1113*)o

The above experiments were performed with perchloric
acid solutions. At low acid concentrations the perchlorate salt
of the iron-phenanthroline complex precipitates upon the add-
ition of either thallium perchlorate or sodium perchlorate.
Therefore in all succeeding experiments nitrates were used in-
stead of perchlorates.

In nitrate solutions when the thallium(III) concentra-
tion was #.74 x lO’B‘E’or larger, the over-all reaction con-
sisted of two concurrent reactions: the dissociation of the

isomplex(first order in complex) and a second order oxidation

reaction analogous to the oxidation reaction noted in the.
perchlorate solutions by George and Irvine.(Fig, 2)
F 2+
d r. (phen) :1
L L

2 Z-
- at m kdiaaiTe(phen)3;} + kaxfie(phen)3+ (T13+)

 

The two paths were found for a variety of concentra-
tions of the complex when the thallium concentration was
1+.71+ x 10'"3 yin-the hydrogen concentration was 1.0 fl and the

ionic strength was 3.0.

FIGURE 1

fifth-all? ' .
Time dependence of ln-—--z- in perchlorate solution.

in»)

-6.B
-7.0
-‘7.2
-7.#
~7.6,_
-‘7.8
.8.0

m” )

['6 (phon)2*]

1n

 

-8.2

-8 O L"_____ . _ J n A
O 1200 2400 3600 4800

 

Time (sec)

fl"e(phen)§":]0 s 1+.31 x 10-5 .11
(1'13”) . 3.90 x lo"2 5
(3”) - 2.7 g

P 3.0

Temperature 8 25.0 1 o.1°c.

TABLE 1

[T1(III)N0;X m M71} x 10'3 .11 Temperature at 25.0 I 0.100.

(H+) . 1.0 E

/}L- 3.0

 

 

 

 

 

 

[Fe (phenfiifl~ (fl) tom (l/moleZsec) deviation
1.21 x 10‘“ 5.04 x 10’2 0.25 x 10"2
6.00 x 10"5 4.69 x 10'2 0.10 x 10'2
5.75 x 10'5 4.89 x 10'2 0.10 x 10"2
5.60 x 10'5 4.52 x 10'2 0.17 x 10"2
5.30 x 10’5 4.56 x 10'2 0.13 x 10'2
3.22 x 10‘5 4.96 x 10'2 0.17 x 10'2
3.12 x 10'5 4.90 x 10'2 0.11 x 10'2
3.12 x 10"5 4.93 x 10’2 0.14 x 10'2
2.69 x 10'5 4.71 x 10'2 0.08 x 10'2
2.69 x 10"5 4.65 x 10'2 0.14 x 10'2
average kox s #.785 x 10.2 l/mole/sec

standard deviation 2 0.146

x 10.2 1/m01e/sec

 

 

FIGURE 2

Ef‘e (phen)§j .

Time dependence of 1n in nitrate solution.

 

 

 

(1'13”)
-3.66
-3.68p
~3.70
r: '3e72 _
N
E I: 6.71.
a a '3076
N; v
Lt; '3e78
.5 -3.80
‘3e82P
'3e8‘5‘
“3086 . a i
o 200 400 600 800 1000

Time(sec)

@‘e(phen)‘§j0 m 1.21 x 104* g

(113*)0 - 4.74 x 10'”3 g
(3“) - 1.0 g

,u- 3.0

Temperature m 25.0 I O.1°C.

. .
The concentrations have been corrected for the dissociation

of the iron complex.

 

'23.
The rate constant for the reaction at the conditions

listed in Table 1 is 4.79 e 0.15 x 10'2 liter .mole'l sec'l.
Table 2 gives the corresponding data for the reaction at the

same concentrations at 30°C.

TABLE 2
[T1(H03)31 s 1+.7'+ x 10-3.33 Temperature a 30.0 1 0.100.

[3“] - 1.0 g

 

 

 

 

 

 

[vb - 3.0
Fe (phen )3" (.11) It“ (1/mole .sec) . 2
deviatlon (deviation)

5.75 x 10.5 0.515 I 0.13 0.0169
3.16 x 10"5 0.562 0.34 0.1156
2.68 x 10"5 0.524 0.04 ' 0.0016
2.53 x 10"5 0.536 0.08 0.0064
2.16 x 10'5 0.502 0.26 0.0676
average tax I 0.528 l/moleOsec
standard deviation s I 0.17h l/mole-sec

 

 

The rate constant for the above reaction is 0.528 t 0.17h
liter mole”1 sec'l.

When the thallium concentration varied below h.74 x 10'}
moles per liter, the rate of dissociation of the iron complex
increased rapidly as the thallium(III) concentration decreased.

The data suggest that the thallium(III) acted as a catalyst for

the dissociation of the iron-phenanthroline complex. The

solutions which contained thallium(III) nitrate in the ~

25+...

concentration range 0f #.22 x 10.“ 5 to 2.11 x 10-6 .1! did not
turn blue during the course of the reaction, which indicated
that no tris-(l,10-phenanthroline)iron(III) ions were formed.
At these low thallium(III) concentrations the rate increased
very rapidly with a decrease in the thallium(III) concentration.

These results seemed to indicate that the dissociation
reaction rate was much greater than the rate of the oxidation
process. A series of experiments using trivalent ions which
would not oxidize the iron-phenanthroline complex were performed
to obtain data on the dissociation rate which would assist in
the interpretation of the data obtained for the two concurrent

reactions. The data obtained using A1(III), Ce(III) and La(III)

are tabulated in Table 3.

 

 

TABLE 3
[?e(phen)§€] - 6.9 x 10-?! Temperature 2 25.0 1 0.100.
(3*) - 1.0 g '
09*) s 4.7‘t x 10.“ 2!
Ala 3.0 _
Ion need 1 Rate (mole/ liter/ sec)
: experimental-inmuang
A13I . 7.08 x 10"8
Ce3+ 6.88 x 10"8
1..“ _ 7.14 x 10"8

 

 

 

 

25¢~

These rates are analogous to those found for the
thallium(III)-tris-(1,10-phenanthroline)iron(II)"nitrate
kinetic solutions. An experiment was then performed to determine
whether strontium(II) behaved in a similar manner. As the
concentration of strontium(II)increased, the rate of the
disappearance of the iron-phenanthroline complex increased.
(Table A). Further studies on the lanthanum(III)-iron-
phenanthroline system showed that the rate of disappearance
of the iron complex increased as the 1anthanum(III) concen-

tration decreased (Table 5).

 

 

TABLE 0

{Te(phen)§I] I 5.75 x 10"5 3| Temperature s 25.0 : O.l°C.
(11*) - 1.0 g

I 3.0

(Sr2+) Q!) Rate (mole/ liter/ sec)

experimental;;9;3;:l______;m

2.37 x 10'“ 7.0 x 10-8

1.19 x 10"3 9.42 x 10’8

4.74 x 10‘”3 . 9.17 x 10"8

 

 

 

 

TABLE 5

(H+) - 1.0 3.

Temperature s 25.0 t 0.1 g

 

 

 

 

 

 

 

[ALI 3.0

fie<phen)§*] (Lay) Initial Rate 1:1 I 1:2
(3 x 10,5) (3) (dole/liter/sec) (mole/ (sec-1)

liter/
sec)

5.75 4,74 x lo"3 :. 1.06 x 10’7 0.0877 0.0595
2.78 4.74 x 10"3 4.80 x 10'8 0.0885 0.0595
1.67 4.74 x 10'3 2.28 x 10‘8 0.0855 0.0595
7.00 9.48 x 10‘“ 9.00 x 10‘"8 0.0815 0.0595
6.90 4.74 x 10'“ 7.14 x 10'8 0.0733 0.0595
6.10 4.74 x 10'” 8.07 x 10'8 0.0866 0.0595
2.85 4.74 x 10‘“ 3.97 x 10"8 0.0845 0.0595
2.22 4.74 x 10’“ 5.38 x 10"8 0.1015 0.0595
7.00 9.48 x 10'5 9.00 x 10"8 0.0850 0.0595
7.20 4.74 x 10‘5 1.04 x 10" 0.0790 0.0595
6.95 4.74 x 10'5. 1.94 x 10'7 0.0990 0.0595
3.45 ‘ 4.74 x 10"5 8.50 x 10'8 0.1013 0.0595
1.98 4.74 x 10"5 3.71 x 10"8 - 0.1000 0.0595

average k1 s 0.0887 mole/liter/sec

standard deviation - 0.008% mole/liter/sec

 

 

 

27.

The rate constant were calculated from the equation

d (Fe (phen )gfl
- dt '

 

m k1 + k2 [Fe(phen)§"]
which in integral form (see Appendix 1) is

fie(phen)§qt s - kl/k2 +{E‘e(phen)§;]o + kl/kz-gg'kzt.

The values for k1 and k2 are 0.088? g 0.008“ mole liter-lose.1
and 0.0595 sec'l, respectively.

A comparison of the rates for reactions in which low
concentrations of thallium(III) are used with rates for reac-
tions where lanthanum(III) was utilized (see Appendix 3) showed
a good correlation between the two sets of rates. The thallium
could be catalyzing the dissociation of the iron complex, but
the hydrolysis of the thallium(III) probably complicates the
reaction at low thallium and hydrogen ion concentrations and
makes the mathematical interpretation of the results uncertain.

It appears that the thallium catalyzed dissociation may
also consist of two concurrent steps; one first order in iron
complex and the second zero order.

The reaction rate for both the thallium(III) and the lan-
thanum(III) reactions with tris-(l,10-phenanthroline)ir0n(II)
ion were accelerated as the hydr0gen ion concentration was

increased at high acidities. When the acid concentration was

\II III III 1.! II I ll
8 II

28.
lowered below 0.7'fl, the rate appeared to increase with decreas-

ing hydrogen ion concentration.

.qei' "

TABLE 6

[Ee(phen)§"]- 6.00 x 10"5 32

(T13*) m 4.74 x 10"3 g
(1.9“) m 4.74 x 10"3 5

Temperature = 25.0 : 0.10C.

 

 

 

 

 

,0,- 3.0
(If) (11) Initial Rate, experimental
(mole/liter/sec)
T1(III) La(III)

2.9 2.69 x 10'7
2.7 1.96 x 10'? ‘
2.5 2.07 x 10'7 E
2.0 1.34 x 10'7 5.56 x 10’7 E
1.5 7.87 x 10‘”8 1.39 x 10"7 §
1.3 1.04 x 10'7 i
1.2 6.57 x 10'8
1.0 1.86 x 10"8 7.14 x 10'8 i
0.7 2.58 x 10’8
0.5 9.56 x 10'9 8.84 x 10"8

 

 

3.9.;

The increase in the rate of disappearance of the iron-
phenanthroline complex with the increase in acid strength was
much greater in the thallium case than in the lanthanum case.
Since the tris-(l,lO-phenanthroline)iron(II) ion concentration
was equal in both cases, the increase could be due to the oxi-
dation of the iron complex by thallium. The increase in the
oxidation rate with an increase in hydrogen ion concentration
might be caused by a smaller degree of hydrolysis of the
thallium(III) ion—- the oxidation probably involves the T13+
species. 8

(George and Irvine(13) reported that the rate for the
thallium(III)-tris-(1,10-phenanthroline)iron(II) ion was
approximately inversely proportional to the first power of the
hydrogen ion concentration. This relationship is not found in
the present work. In reactions where there was proof of the
oxidation of the complex, the rate increased as the hydrogen
ion concentration increased,

The above result serves to cast doubt on the assumption
of George and Irvine that the chief species reacting with the
complex ion reducing agent is T10+. At.1ow acid concentrations
when the disappearance of the iron complex was complete, it was
noted that there was still T1(III) present in the solution. The

T13+ concentration may be too small to oxidize the iron complex,
and/or the low T13+ concentration may catalyse the dissociation

reaction greatly.

M

31....__

A study of the effect of various univalent ions was
performed(5) by adding varied amounts of hydrogen ion,
potassium ion, lithium ion and nitrate ion to a solution of
tris-(l,10-phenanthroline)iron(II) ion which was 2.5 x 10"5 a.
A simultaneous increase in the hydrogen and nitrate concen-
trations caused an increase in the rate of dissociation. The
same result was obtained with lithium nitrate, but the addition
of larger amounts of potassium nitrate caused a slight increase
in the rate. These rate variations were small compared to those
noted for the addition of 1anthanum(III) or thallium(III)
nitrate to the iron-phenanthroline complex. This work done by.
Dickens, Basolo and Neumann gave a rate constant of approxi-
mately 7.5 x 10'"5 sec-1 for the iron complex dissociation.

The present experiments that were performed by using
perchlorate solutions give results analogous to the nitrate
date. The perchlorate results must be weighted rather lightly
since precipitation appeared to occur during the kinetic study
of most of the solutions. Only in the 2.7 E acid solution
precipitation definitely did not occur.(See Appendix 8)

These experiments except for the hydrogen ion dependence
seem to support the work of George and Irvine when large con-
centrations of thallium(III) are used, and it can be expected
that the thallium(III) concentration will be nearly constant
throughout the experiment. Under these conditions oxidation of

the tris-(l,10-phenanthroline)iron(II) ion appears to occur,

32.2.
but dissociation of the complex seems to occur concurrently.
This latter supposition was not considered by George and Irvine.

Atlow thallium(III) concentrations there appears to be
little or no oxidation. The principal reaction seems to be a

rapid dissociation of the complex which is catalysed by the

unhydrolysed trivalent ion.

1.

2.
3.

#.

5.

9.
10.

11.

12.

13.

14.

15.

16.

17.
18.

BIBLIOGRAPHY

Williams, R. J. p., J. Chem. Soc., 137 (1955).

Kolthoff, I. N., Lee, T. S. and Leussing, D. L., Anal. C em.,
22. 985 (1948).

Basolo, P., Hayes, J. and Neumann, H. H., J. Am. Chem. Soc.,

6, 3807 (1954).

Lee, T. 5., Kolthoff, I. M. and Leussing, D. L., J. Am. Chem.
Soc., 29, 2348 (1948).

Dickens, J. E., Basolo, F. and Neumann, H. H., J. Am. Chem.

Soc., 22, 1286 (1957).

Prestwood, R. J. and Jahl, A. C., J. Am. Chem. :oc.,'Zl,
3137 (1949).

fiarbottle, G. and Dodson, R. J., J. Am. Chem. $00., 22,
anuz (1951).

Eorscheimer, O. L. and Bpple, R. P., J. Am. Chem. Soc., 23,

5773 (1952).
Johnson, C. E., J. Am. Chem. Soc., 23, 959 (1952).

Born, H. J., Vogt, H. and Vogt, G., Z. Tlektrochem., fig,
372 (1962).

Dulz, G. and Sutin, N., Inorg. Chem., 2, 917 (1963).

Campion, R. J., Purdie, N. and Sutin, N., Inorg. Chem., 2,
1091 (196%).

George, P. and Irvine, D. H., J. Chem. Soc., 587 (l95h).

Sutin, N. and Gordon, B. N., J. Am. Chem. Soc., Q}, 70
(1961).

Smith, G. P., McCurdy, w. H. and Diehl, H., Analyst, 21,
418 (1952).

Zintl, E. and Rienacker, G., Z. anorg. allgem. Chem.,

Kolthoff, I. M., Rec. Trav. Chim., 31, 172 (1922).

Sherrill, M. S, and Haas, A. J. Jr., J. Am. Chem. Soc.,
8. 953 (1936).

.35}.

.
8

APPENDIX 1

Let [Fe(phen)§+] a y ; then - 3% :- k1 + kzy

t
d! :2 2 . - 1:1.th

dt

t t
yek2 - «(kl/kahk2 + c

«at
When t . o, c . yo + (kl/k2).
.th

1"kl/k2+(yo+kl/k2).

 

APPENDIX 2

 

 

(8*) - 1.0 g /u.- 5.0 Temperature . 25.0 t 0.1%.
{hemogfl (g) (113*) (5) (1.3“) (13,) mm: mm,"
experimental
(molellgterZIth ‘
5.57 x 10‘5 _ 2.11 x 10"“ 1.95 x 10"7
6.16 x 10’5 ' . 4.22 a 10'5 1.21 x 10'7
6.00 x 10"5 2.11 x 10"5 i 3.2!» x 10"7
6.20 x 10’5 11.22 x 10'5 5.15 x 10”
5.75 x 10'”5 2.11 x 10'6 2.61 x 10:7
7.00 x 10"5 9.#8 x 10'# 9.00 x 10'8
6.90 x 10'5 1571+ x 10'“ 7.1l+ x 10"8
7.00 x 10'5 9.1.8 x 10"5 1.35 x 10"7
7.20 x 10"5 - 1+.7h x 10"5 1.0t+ x 10‘7“

 

 

 

 

 

APPENDIX 3

De(phen)§q I 5.75 x 10"5 g

(133*) a £1.74 x 10'3 g

(3*) . 1.0 fl

’4.- 3.0

 

Temperature ( C. )

Inital Rate, experimental

 

35.0
30.0
20.0

 

 

Mtg-[sea
1.05 x 10'7

9e00 I 10-8

5e 51" X 10.8

 

 

 

(3*) - 0.10 g

./1- 3.0'

APPENDIX 4

Perchlorate Data

Temperature m 25.0 t 0.100.

 

 

 

 

 

[p.(paeaigt‘l (g) mi”) (5) mm Rate, experiments]
‘ (moleZliter/sec)

1.07 x 10"“ 1.07 x 10"“ 3.28 x 10’8

1.6h x 10‘“ l.h6 x 10:z\ 2.67 x 10'?

1.07 x 10'“ 9.70 x 10'3 9.60 x 10'9

2.57 x 10‘5 1.80 x 10"3 3169 x 10"8