THE KINETICS OF THE OXIDATION OF
HYPOPHOSPHITE BY PERRUTHENATE
m AQUEOUS ALKALI

Thesis for the Degree of Ph. D.
MICHIGAN STATE UNIVERSITY
SAMUEL J. PATON
1972

“3 LIBRARY

Michigan State
University

 

This is to certify that the

thesis entitled

THE KINETICS OF THE OXIDATION OF HYPOPHOSPHITE BY
PERRUTHENATE IN AQUEOUS ALKALI

presented by

Samuel J. Paton

has been accepted towards fulfillment
of the requirements for

Eh D degreein Chemistry

Major professor

Date October 13, 1972

0-7639

   

    
 
     
      

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ABSTRACT
THE KINETICS OF THE OXIDATION OF HYPOPHOSPHITE

BY PERRUTHENATE IN AQUEOUS ALKALI

The kinetics of the oxidation of hypophosphite by
perruthenate was studied at various temperatures in
0.300 M KOH. The change in perruthenate concentration
was followed spectrOphotometrically under pseudo first
order (stOpped-flow) and second order conditions. Data
were analyzed by a curve fitting computer program. The
reaction is first order with respect to each anion and
hydroxide with a rate constant of zue flfzscc'1 at 24.80.
The rate decreases linearly with decreasing pH. The
Arrhenius activation energy E8 is 2.h6 kcal/mole and the
entropy of activation.ASq= = -hl.3 cu. Results are compared
to published studies of hypOphosphorus acid in acidic
solution and hypophosphite anion in-basic solution. The
proposed mechanism involves an equilibrium with hydroxide

ion and hypOphosphite. The activated complex is probably

a hydroxide bridged species of hypOphOSphlte and

perruthenatc.

THE KINETICS OF THE OXIDATION OF HYPOPHOSPHITE
BY PERRUTHENATE IN AQUEOUS ALKALI

By

I. {r

Samuel erpaton

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

DOCTOR OF PHILOSOPHY

Department of Chemistry

1972

 

33 e. .Nflflnfii I...
Tani-Hi
. .

 

ACKNOWLEDGEMENT

The author gratefully acknowledges the patience and
understanding of Professor Carl H. Brubaker, Jr..
financial aid from National Defense Education act (Title IV)
and Dow Chemical Company, and a little help from his

friends.

TABLE OF CONTENTS

IntrOdIICtion o o o o o o o o o o o

EXperimental . . .,. . . . . . . . . . . .

A. Preparation and Analyses of Reagents

E. Second Order Kinetics. . . .
C. Pseudo-First Order Kinetics.
D. Variation in Temperature . .
E. Variation of pH. . . . . . .
F. Products of the Reaction . .
Results. . . . . . . . . . . . . .
A. The Rate Equation. . . . . .
B. Activation Parameters. . . .
C. pH Dependence. . . . . . . .
Discussion . . . . . . . . . . . .
Appendix A . . . . . . . . . . . .

AppendiXBoooooooooooo

ii

XV

 

III

IV

VI
VII

VIII

IX
A - I
A- II
A-III
A- IV
A - V
A- VI
A-VII

LIST OF TABLES

Half-Life Values for Ruou' Reduction

The Rate in 0.300 M KOH at 2u.8° . .

The Rate in 0.300 m KOH at 2n.8°

Modified . . . . . . . . . . . . .
The Rate in 0.300 M KOH at u.3° .
The Rate in 0.300 M KOH at 14.7° .
The Rate in 0.300 M KOH at 34.3° .
The Rate in 0.300 M KOH at u2.8° .

Depepdence of the Rate on [OH-J at
2u08 C O C O C C C C O O O O O C 0

Representative St0pped - Flow Data
Absorbance at lh.7°C, Experiment 1
Total Ru Concentration. Experiment
Absorbance at lb.7°C, EXperiment 2
Total Ru Concentration, Experiment
Absorbance at lh.7°C. EXperiment 3
Total Ru Concentration, EXperiment

Extrapolated RuO ' Concentration
att=0.0seC..........

iii

21
25
26
27
28

32
3?

vii
vii
viii

viii

ix

Figure

II

III

IV

A- II

A-III

LIST OF FIGURES

Extent of Reaction gs Time, 21+.8o .

Extent of Reaction gs Time, 24.8o ,
2 Parameters. . . . . . . . . . . .

Plot of Arrhenium Activation Energy
Function. . . . . . . . . . . . . .

Dependence of the Rate of [OH-] . .

Extent of Reaction gs Time,
Hypothetical. o o o c o o o o o o o

O
Extent of Reaction gs Time, l#.7 .

Extent of Reaction gg Time, 1a.7° ,
2 Parameters. . . . . . . . . . . .

iv

23

30

3h

iii

xii

xiv

INTRODUCTION

P. L. Dulong, in 1816, treated phosphides of alkaline
earths with water, isolated hyponhosnhite salts, and
proposed the name hypOphosphorous acid.1 Since then the
mono protic acid and many of its salts have been studied.
Detailed studies of the kinetics of reactions involving
HBPO2 as a reducing agent began with Mitchell.2 He
suggested an acid catalyzed equilibrium between HBPO2 and

an active form of the acid

H3P02(inactive) + H++>H3P02(active) + H+

which was immediately oxidized.
K

II
H3P02(active) + 0X -----'>H3P03 + (reduced product)

The general form of the rate law,

 

_ d(H3P02) z k(H3P02)(H+)
dt 1+LL1£L

N

k (0X)

is obeyed for the following oxidants: 12, Brz, 012, 103',
HCrOnI, Ce(IV), and Tl(III) (References 2 through 7,
respectively). Evidence for complex formation between
HBPO2 and the oxidant was presented in the latter two

studies. Isotope exchange studies by Jenkins and Yost8

(T/H) and Fratiello and Anderson9 (D/H) further substantiated

I

2
the inactive/active phenomena for H3P02' Other rate laws
(involving complex formation as well as the inactive/active
concept) were observed by Carroll and Thomas10 (Ce(IV)),
Cooper et. a1.11 (V(V)), and Cooper12 (Cr(VI)). All of these
studies were in acidic or neutral solutions. Studies of
reactions of hypOphOSphite in basic solutions are limited.
In 1912, Sieverts and Loessner13 studied the rate of the
decomposition reaction at 910 and 1000 in 1-4 M NaOH
solution. Their data conformed to the rate equation

' dt

 

= k(H2P02')(0H‘)2

and they observed a "normal temperature dependence".
Jenkins.and Yost8 mentioned decomposition problems in

their work at high pH. Roper, Haas and Gillmanll+ published
results of a proton nmr study on the base catalyzed
hydrogen-deuterium exchange on hypophOSphite anion
(discussed in detail under RESULTS). The authors suggest
the rate of production of the reactive intermediate was

.15

observed. Ben-Zv1 studied a free radical chain mechanism
of oxidation by peroxydisulphate up to pH=1l. No other

basic solution studies have been published.

This study was intended to observe effects of a
reactive intermediate of hypOphosphite in basic solution.
It was necessary to use an oxidant known to gain electrons
rapidly. Luoma16 demonstrated the extremely rapid exchange
between Ruou- and RuOu= as well as a fast reaction between
Ruou- and Mn0u=. Thus, the choice of KRuO“ as the oxidizing
agent is reasonable.

EXPERIMENTAL

A. PREPARATION AND ANALYSIS OF REAGENTS KHZPOZ

 

About 400 g of Baker and Adamson KHZPO2 was added to
300 ml of distilled water, and heated to 50) to facilitate
dissolution. The warm solution darkened. The grey matter
was removed by filtering through medium porosity filter
paper. The clear solution was cooled to 70 in a refrigerator
but no crystals appeared. The solution was transfered to a
Rotovap ® and maintained at 300 by means of a water bath
and the apparatus was evacuated to about 10 torr. Crystals
of KHZPO2 formed suddenly and were removed by filtration on
a coarse frit. After several washings with distilled water,
the white crystals (200 g) were stored in a brown bottle.

The supernatent was returned to the Rotovap ® evaporator for

14.
second (50 g) and third (20 g) batches of crystals. The
three samples of KHZPO2 were dried in vacuo over Drierite ®
for several hours, then eXposed to the air overnight. Two
samples from each of the three crystallizations were

17 The determined

analyzed by the method of Ogawa.
molecular weight ranged from 104.50 to 105.93 with an
average of 105.05 g/mole (actual molecular weight = 104.09
g/mole). The three batches of the salt were crushed and
thoroughly mixed. Four samples of about 0.1 g each were
analyzed by the two part method of Jones and Swift.18

The first part, a sensitive analysis for phosphite (the

most likely contaminant), showed none was present. From

the second part, the oxidation of all P(I) and P(III)
compounds to phosphate. the determined molecular weight was
105.07 1 10 g/mole. Galbraith Laboratories found 29.50% P
which correSponds to 105.0 g/mole. Actual %P for KHZPO2

is 29.76%. The value of 105.06 for the molecular weight

of KHZPO2 was used for all future calculations. The 0.9%
difference was attributed to water adsorbed on the surface
of the crystals. The assumed molecular weight was confirmed

several times by the method of Jones and Swift during the

course of the study.

KRUOu

 

Ruthenium metal powder (99.9%) from K&K Laboratories
and anhydrous RuCl3 (99.9%) from Alfa Inorganics were the
sources of ruthenium. The starting material KRuOu was
prepared in the following manner. Tan g of sodium ‘
peroxide and three g of Ru powder were fused in a nickel
crucible over a Meeker burner. After cooling, the fused
mass was dissolved in 300 m1 of distilled water. The
solution was carefully acidified with 6 M HCl, which caused
the orange solution (mostly Ru04=) to turn black (due to the
formation of hydrated Ru(IV) oxide). The pH was adjusted
to 7 with NaHCOB. Hydrated oxide was removed by filtration
on a medium frit. dried at 110°, and ground into a powder.
(After each experiment Ru was precipitated as the oxide and
recycled at this point.) The black powder was evaporated
to dryness with concentrated HCl. The water soluble RuCl3
product (which contained a small amount of the hydrated
oxide) was the starting material for the preparation of
RuOu. The volatile tetraoxide was generated according
to the recipe of Larsen and Ross.19 A typical preparation

was carried out in a one liter, three neck (zu/uo) round

6
bottom flask. An overhead stirring assembly was placed in
the central neck. A 7 mm glass tube for N2 flow in a 2h/25
Teflon ® adapter (RuO,+ reacts quickly with all st0pcock
grease, and even discolors Teflon CD was placed in another
neck. The third neck contained another 24/25 adapter
which had a U-shaped tube (10 x 25 x 10 cm). This tube
delivered Ruou (under slightly positive pressure of N2)
into a 100 m1 flask equipped with a side arm. The receiving
vessel was submerged in an isoprOpanol/dry-ice bath.
Reagents used were 3.5 g of “RuClB”. 125 g NaBiOB. 500 ml
6N H280“, and 200 ml of distilled water. As the mixture was
stirred, N2 was bubbled through at 3 to 5 bubbles per
second. Heat was applied from a heating mantle to bring
the mixture to a temperature near 90°. By this time bright
yellow RuOu (as well as ice) was observed in the receiving
flask. When RuOu no longer condensed in the delivery tube.
the reaction was judged completed. The contents of the
receiving vessel were quickly added to an equal volume
of 0.3 M_KOH solution (total volume 80 ml) in a large
weighing bottle. The t0p was replaced and the weighing

bottle placed in a refrigerator at 70 for 3 days. At 12

7

hour intervals. 3 ml of saturated KC] solutirr were added.
The shiny black crystals of KRuOu were removed from the
orange supernatant on a coarse filter frit. After being
washed four times with small amounts of distilled water the
KRuOu (.3 g) was dried and stored over Drierite ® .
(Calculated for KRuOu: Ru, h9.50. Found = 49.77) The actual
Ruou' and Ruou= concentrations were calculated from
absorbences at #65 and 383.5 nm by use of molar absorptivities
given by Luoma.16 (Total Ru concentration was consistently
2-3% lower than the value calculated from the weighed amount of
KRuOu.)
£921 ‘

The reaction medium of 0.3 M KOH was prepared fresh
for each experiment from Baker's Analyzed Reagent CD which
contained less than 0.8% carbonate. Sufficient pellets to
give slightly more than a 0.3 M solution were dissolved in
distilled water (deionized water distilled from KMnOu and
then from itself). From this solution exactly 100 ml
were removed and used to titrate weighed samples of potassium
hydrogen phthalate to the phenolphthalien endpoint.

Distilled water was added to dilute the stock to exactly

0.300 M,in KOH. All other reagents met ACS standards and

were used as taken from their containers.

B. SECOND ORDER KINETICS

When Ruou- and HZPOZI concentrations were similar and
near 2 x 10'“ molar, the reaction time was of the order of
minutes. A Cary Model 1h spectrOphotometer was used for
monitoring the concentration of Ruou' or Ru04= as a function
of time. A typical experiment began with the calibration
of 5 cm cells made of Pyrex (silica seemed to affect the
stability of Ruou') at both #65 nm and 383.5 nm (the
adsorption maxima for RuOu= and Ruou', respectively).
A stock solution of 0.300 M KOH was prepared as described
previously. A 0.131 g sample of KHZPO2 was dissolved in
the hydroxide and the solution diluted to 250 ml. From
this solution, 25.0 ml was pipetted into another 250 m1
volumetric flask and again diluted to the mark with KOH.
The resulting solution was 5.00 x 10-“ M’in KHZPOZ'

A 500 ml volumetric flask was partially filled with
300 ml of the stock 0.300 M KOH. The stock solution of
KRuOu was prepared by weighing 0.0h51 g of KRuOu into the

flask and quickly swirling the contents. (It was found

9

that KRuOu crystals caught along the walls and rapidly
decomposed to Ru04= unless quickly brought into solution.)
The resulting solution was u.u2 x 10”+ M in KRuOu.

The three stock solutions (KOH, KRuOu, and KHZPOZ)
were stored in a controlled temperature water bath. (Water
from this heat sink was pumped through the cooling passages
of the cell compartments of the Cary 14). Aliquots (25.0 ml)
of the hypophosphite and perruthenate were pipetted into
separate 100 ml beakers which were held in the bath. The
beakers were removed and the perruthenate solution was
quickly poured into the hypOphOSphite solution. (This
action was t = 0.0 sec and a timer was started.) The
mixture was poured between the beakers four times, the
sample cell was rinsed twice, filled. and placed in the
sample chamber. The recording pen was actuated immediately
(about t = 25 sec) and the chart paper started at
t = 30 sec. The absorbance was recorded at #65 nm until
at least 90% of the Ruou' was converted to Ru04=.
Absorbance readings at 383.5 nm were recorded at several
times during the course of the reaction. The experiment
was repeated several times alternating hydroxide with

hypOphOSphite.

TREATMENT OF DATA

Absorbance was read from the chart at minimum of 15
different times which were selected such that half were
in the first quarter of the reaction (whereadA/At was
greatest). An average value for total Ruon- + RuOu=
concentration was determined from the absorbances at
383.5 and 465 mn. With this value and an absorbance, the
perruthenate concentration at any time “t" could be
calculated. A small correction factor (20) to account for
oxidizable contaminants in the KOH pellets was applied to
the value of RuOu' concentration. The value thus obtained
represented the concentration of Ru0“' that had not yet
reacted with HZPOZI. This variable was fitted to a rate

law as described under “Results".

C. PSEUDO FIRST ORDER KINETICS

Two different stOpped-flow apparatus made it possible
to study the reaction over a wide range of concentrations.
One of the home-built instruments delivered about 15 ml of
each reactant solution into a mixing chamber. The flow
continued through 2 mm capillary tubing to a solenoid

stop valve and then to a catch vessel. A beam of

IO

11
light from a monochomator passed through the 2 mm tubing
perpendicular to the flow of liquid and then into a
photomultiplier. The amplified output voltage (directly
proportional to %T) was logarithmically converted and
stored on a recording oscilloscOpe. Voltage (now pr0portional
to absorbance) versus time traces were photographed with
Polaroid ® film. Treatment of these data is described
later.

The other st0pped-flow instrument21

was basically
similar except for light path length. Each syringe pushed
about 1 ml through a mixing chamber into an observation
cell of 2 mm stainless steel tubing 2 cm long. Light
traversed the length of the cell, thus giving a path length
10 times the system previously described. Although the
two instruments exhibited major detailed differences, the
principles of operation were identical.

In a typical eXperiment (on either instrument), the
oscilloscope recorded a trace representing zero absorbance
(0.300 M KOH in each syringe) at #65 nm. The RuOn'

solution was admitted to one syringe and the trace

corresponding to the absorbance of RuOu' only was recorded.

12
These traces were photographed, and were the basis for all
absorbance assignments. The other syringe was loaded with
HZPOZI and the reaction trace was recorded and photographed.
After a thorough rinsing, the entire procedure was repeated
with solutions of different concentrations. Several
reactions (traces) were recorded and photographed at each

concentration.

TREATMENT OF DATA

The grid pattern of the photograph of the oscilloSCOpe
was scaled for absorbance and time. Absorbance was con-
verted to RuOu- concentration. The absorbance value at
t = 0.0 see was scaled to 1.00 and all other concentrations
divided by that value. Concentration versus time was
plotted on semi-log paper, and the half—life for each

reaction estimated from the graph.

D. VARIATION IN TEMPERATURE
A 100 L water bath equipped with circulating pump,
tap water cooling coils, and thermostatically controlled
heating blades provided a constant temperature of 21+.8o
for the cell compartment of the Cary lb. The temperature of

the solution in the 5 cm cell in the sample compartment

13

was always within 0.1° of that value. A 50 l.coolant bath
unit manufactured by the Wilkens-Anderson Company was used
to control temperature above or below ambient conditions.
A thermometer was inserted into the 5 cm cell at the end
of each reaction to measure the actual temperature. Data

were treated as described previously.

E. VARIATION OngH

Potassium hydroxide pellets and K3POu were mixed in
various proportions to vary the pH while maintaining the
ionic strength (on a molar scale) at 0.30. A Corning
Model 10 pH meter equipped with a general purpose glass
electrode and a new referance electrode indicated the
actual pH of the various solutions. Measurements before and
after the reaction differed by less than 0.1 pH unit. Data

were treated as described previously.

F. PRODUCTS OF THE REACTION

It was clear that perruthenate was converted to

ruthenate and that the rate of reduction of RuOu— by

Hzpo; was not significant at these concentrations and time

scales. Indeed, knowledge of the total RuOu‘ + RuOL:=

concentration was an integral part of the data analysis.

In

Inconclusive results were obtained from the several
attempts to establish the nature of the product of the
oxidation of hypOphosphite. A 50 m1 aliquot of 8 x 10"3
M RuOu' was allowed to react with a 5 ml aliquot of
5 x 10"2 M HZPOZI. When all ruthenium was in the +-6
state. the solution was carefully neutralized with 6 M HCl.
A small amount of saturated NaHCO3 was added to adjust the
pH. Hydrated Ru(IV) oxide precipitated as the solution
approached neutrality. The solution was centrifuged, and
the decanted portion filtered through a medium frit. The
phOSphite and hypOphosphite were determined in the very
pale yellow solution by the method of Jones and Swift.l8

Thin layer chromatography identified the only reaction
product to be phosphite. The plates were Baker-Flex brand
#0-h468 (cellulose). Seiler'szz recommendation for the
solvent, a mixture of methanol concd ammonia-10% trichloro-
acetic acid-water (50+15+5+30), gave excellent results.
Substitution of sodium molybdate for ammonium molybdate
caused no difficulty. The Rf value for HP03= (compared

to H P02-) was 0.77 for reference and 0.76 for sample.

2

15
No paramagnetic species were observed in the frozen
(liquid N2) reaction solution. Luoma16 also failed to
observe an ear signal from RuOu-. The samples were run

on a Varian model E-h at liquid N2 temperature and were

scanned from 1500 to 6500 Gauss at 9.295 GHZ.

RESULTS

A. THE RATE EQUATION

Extensive stepped-flow studies were designed to
observe the effects of an equilibrium between inactive and
active forms of H2P02-. Table I presents the half-lives
for the disappearance of RuOu' under various psuedo first
order conditions. The t% values are consistent and all
plots are linear for at least two half-lives.

Data from dilute solution studies were analyzed by a
curve fitting computer program written by Dye and Nicely.23
The second order equation

dx

E = k2(Ru - x)(P - 325) (1)

16
Table I

Half-Life Values for RuOu- Disappearance
Initial Concentration (M x 10 )

RuOu

 

 

[E332§:] t$(sec)
1000. .080* .089 .10 .13

500. .18 .17

250. .355 .39 .40 .QO .40 .48

125. .81 .88
50. 2.1 2.2
25. 0.25 0.3
12.5 8.3 8.7

5.0 20.5 19.

* All values in sec. 1 10%

17
was integrated and rearranged to the form

ykzt
- -2
x = 2Ru(e l) y = 325.2 (2)
Ru yth

fi— e -2
where Ru and P represent initial concentrations of RuOu'
and HZPOZI, k2 is the second order rate constant, t is
time, and x is the extent of reaction (= the concentration
of RuOh' converted to Ru0n= at any time t). The input
data are x and time, and k is Optimized to give minimum
residuals between the input and the calculated x.

Results of these calculations appear in Table II.
Figure I is a computer plot of the input variable and the
u M

calculated x Mg time for KHZPO = 2.5 x 10'

2
EXperimental points (input) are designated with x, 0
indicates a calculated point, and = means experimental and
calculated points differ by less than the experimental
uncertainty (usually less than 5%).

The perruthenate concentration can be introduced as
a second adjustable parameter in equation (2). The curve-
fitting program will then optimize both parameters. Such
calculations are listed in Table III and a representative

u M

computer plot is Figure II (for HZPOZI = 2.5 x 10' _).

Table II

The Rate of the Reaction in 0.300 M KOH at 2#.8°C

[Hzpoz

1.50
1.00

0.500

Initial Concentrations (M'x 10h)

[RuOu-]

1.72

 

1.66

|+ |+ [4- H-

H-

19

O

m.:m

.oEHB MN Cowpomom ho peopxm

H musmfim

20

own H opsmwa Aoomv mafia o:

 

 

 

amOIIImlIIOVIIllmIlllmiltlmIIIImIIIImIIIImltllmtIllmllllmltltm | m m u m I mIUIIWIIIOVllllu
m xm
~ OH
n

n

u

m

a I
~

—

H

m K
~ 0
u

u

a

m I

n

u

n

n I

m

u x

n 0

~

u x

m 0

~

~ u

~

u 0

m x

a I

n

n I

u

T I

u C

a o x

a x

u N

m

a

a O

m K

~ 0

m K

a O

n x

ml

n

umIIIImIIICmIIIImOIltmllIIWIIOIWIIIIWIIIImIIIImoIllMlIllm llm M m Im M MI Ml w I

 

Hue-ouch-mnunummmnvdmwuuumI-n—IwmuHunmnnuum—unumuumnmuuu

Imam.

OH x x

Ion.a

[H PO '

2 2

2.50

2.00

1.50

1.00

0.500

]

The Rate in 0.300 M KOH at 24.8°0

21

Table III

Initial Concentrations (M x 10h)

Lied
1.72
1.73
1.69

1.69

1.66

calculated

[R“°u-]

1.76 +

 

1.71 1
1.73 i
1.77 1

3.28

H-

.01
.01
.01

.01

Avg*

 

*Exclude last value

H-

H-

|+

.3
.5

22

meoposmpmm N

.oEHe MM coflpoMom mo enovxm

HH oeamaa

213

‘ .."‘ u'c'
y”:|

com HH ohsmfim Aoomv mews o:

 

 

.m w w m u----m o w muuuuma-c-mu---m--uum----m----m-u--m----m----m--u-m--unmuuuna
m am
a a
_ a
_ a
a a
m m
_ u a
a a
_ a
a _
m u m
_ a
a _
a _
a x _
m o m
a a
a a
_ u a
a a
m m
_ x a
_ o a
a a.
_ u a
m m
a a
a n a
_ a
a. a a
m m
_ u a
a a
e a a
a a
m u m
a O u
_ x a
a u a
a u a
m m
_ a
a a
a u a
_ a
m u m
e _
a u u
ax _
.o _
amuutlm-t--mnu--m----m-tutm--u-mu--umun--m--u-mu.--m---um m m m m m p m m m u a

 

umam.

Ava on x

IHh.H

2“

It was judged desirable to minimize machine manipulations
of the data, so all results are calculated from equation (2)

with k2 the only parameter.

B. ACTIVATION PARAMETERS

All temperature studies were carried out in 0.300 M KOH.
Results from these reactions are listed in Tables IV, V,
VI, and VII. The Arrhenius activation energy Ea’ calculated

from = Ae-Eg/RT

k3
is 2.“5 1 0.2 kcal/mole and the pre-exponential factor

A = 1,63;x105M- sec-1 (see Figure III). Activation

enthalpy (AHIU and entrOpy (ASTW are found from the equation

where As‘F _ AH;
k2 = %I e R e RT
kr = rate of reaction
k = Boltzman's constant
h = Planck's constant
R = gas constant
T = temperature
AH‘F = E - RT
a
The activation parameters are ART: = 1.86 1 02 kcal/mole

and As:F = -41.3 1 7 cu.

25

Table IV

The Rate of the Reaction in 0.300 M KOH at u.3°

Initial Concentrations (M x 10“)

 

 

k2
M2302:] [RuOu’] Mfl sec"1
2.09 1.70 43.9 I .n
1.25 1.77 45.5 i .6
2.u9 1.77 #5.3 1 .3
1.25 1.71 u3.3 1 1.2
Wino

26

Table V

The Rate of the Reaction in 0.300 M KOH at 1n.7°

Initial Concentrations (M x 10“)

[RuOu-1

1.81

 

1.79

1.79

1.80

 

2
11-1 sec"1
148.5 1 .3
5n.1 1 .6
49.1 1 .u
“5.1 1 .6
#9.2 + 1.0

27

Table VI

0

The Rate of the Reaction in 0.300 M KOH at 34.3

Initial Concentrations (M x 10“)

 

 

2
[Mg-1:03;] [RuOu’] M’l sec"1
2.53 1.57 ou.1 1 .7
1.90 1.56 70.0 _+_ .7
1.27 1.50 73.1 1 1.0
0.633 1.t+7 66.3 1 1.7
71.7; 1 2.0

28

Table VII

The Rate of the Reaction in 0.300 M KOH at 42.80

Initial Concentrations (M x 10“)

HZPOZ‘] [RuOu-l .m'l sec'l

 

2.53 1.65 73.7 .1
1.90 1.65 66.3 1
1.27 1.61 84.9 1
2.53 1.45 73-6 1

71:6:

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_2V

C. pH DEPENDENCE

A positive hydroxide ion dependence for the rate
constant was expected. Experiments showed a large effect,
and therefore the reaction was studied over a wide range
of hydroxide concentration. At the lowest pH (10.3), the
reaction solution was turbid when removed from the Cary 14.
It was obvious other reactions interfered at lower pH.
Table VIII presents rate constant results from pH studies.

A plot of k y§_(OH') has a sIOpe of 248 (See Figure IV),

2

which indicate" a positive OH- concentration dependence.

DISCUSSION
. , 14
R0per, Haas, and G111man 8 rate data for the
exchange reaction in D20
HZPO2 —> HDPO2 —->02PO2
obeyed the rate law
rate = k(00‘)(H2P02') (3)

1

with k = 3 x 10’3.M’lsec‘ at 25°C. They concluded

31

[HZPO2

2.64
3.08
2.58
3.49

2.53

32

Table VIII

Dependence of the Rate on [OHi] at 24.80

Initial Concentration (M x 10 )

1 [non

[on-1

mole/liter.

 

0.0833
0.0158
0.00851
0.00363

0.000200

4

 

16.5
5-15
4.17
3.78

3.00

.l

.09

.03

.07

.05

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35

"....we believe it (the rate controlling step)
to correspond to the removal of a hydrogen
from the hypophosphite anion to foEm a triply
connected phOphorous dianion, HP02 1 viz.,
H2P02‘ + OD‘-—>HP02’ + HDO. This reactive
species then will pick up an additional
hydrogen or deuterium atom from the solvent

in a rapid step: HPOZ ' + DZO-d>HDP02’ + OD'.
If this view is correct, the rate to which
equation (3) refers is actually the rate of
production of the reactive intermediate."

If this reactive intermediate is the Species pr0posed by
hitchell2 (i.e., the form that is readily oxidized), then
it should be easily oxidized by the strong
oxidant RuOu'. If the reactive intermediate concentration
is substantially less than the inactive form (say 10%),
the reaction should clearly deviate from pseudo-first
order kinetics before the initial H2P02' to RuOu' ratio is
below 20:1. The t5 values in Table I do not deviate from
the expected pattern. Therefore, the reactive Species is
at least 10% of (and quite likely equal to) the HZPOZ-
concentration at all times during the reaction.

when the ratio of Hzpoz’ to RuOu- is near unity, the

reaction data conform to the second order eXpression

d .
3% = k2(Ru - x)(P - g)

36
with k = 58.5 m‘1 sec'1 at 25°C. Stopped flow data fitted

to equation (2), treated as described previously,

gives similar values for the second order rate constant
k2 (see Table IX). In order to test for a contribution
to the rate eXpression from a OH- concentration term, the

equation

%% = (R2 + k3(0H'))(Ru - X)(P ' I)

was tested. A plot of k2 (the observed k) 11,(08') has
a lepe of 248. 1 14. and an intercept of 0.2 1_1.8.
The hydroxide ion dependence in the rate law must be
first order, and k2 is zero within experimental error.

The overall rate law is

at:

e k3(Ru - x)(P - §)(0H’)

- -1 -
with k3 = 248. 1,14 M 288C when.OH is in excess.

A. G. Miroshnichenko and V. A. Luneok-Burmakinazu

published the results of the oxidation of HBPO2 by H202
in acid solution. Although their conclusions appear self

180 studies) that

contradictory, they do claim (from
peroxide rather than water is the source of oxygen in the

formation of phosphite. In this study, OH- or water must

[H P0“

2 2

1000.
500.
125.

25.

]

37

Table IX

Representative StOpped - Flow Data

Initial Concentration (M x 10h)

calc RuOn]

12.6
6.03
4.56
4.58

4.86

H-

.5
.2
.02

.04

 

\1
O\
C
CD
+
O
\7

 

38

be the source of oxygen, since the exchange of 0 between
H20 and RuOu- is slow compared to the time scale of the
present reaction.26

Ben-Zvils proposed the formation of a hypophOSphite
radical in the oxidation by peroxydisulphate. He listed
the activation energy of the formation of the radical as
6.6 kcal/mole. The rate equation is not simple and the

_ _1
observed rate constant decreased to a plateau (about 10'3 M 2

sec-1) at pH = 11. Although a free radical chain
mechanism for oxidation by RuOu- cannot be excluded, it
could not be substantiated, and the data conform well to
the second order rate law.

R0per, Haas, and Gillmanlh report an Ea of about
19 kcal/mole which is similar to activation energies
from other studies on the oxidation of hypOphosphite.
The low Ea from this study (2.5 kcal/mole) and the large
entr0py of activation (-41 eu) indicates the rate
determining step is other than the inactive/active

tautomerism. The data can be explained by a pre-

equilibrium between H2P02 and OH‘

HPO- + 0H’=Hm'08’

2 2

39
The large, negative activation entrepy indicates
oxidation of H2P02 occurs by means of a bridged
activated complex. The bridge probably involves-OH-
ion. The rate expression indicates one BuO; ion is

involved in the rate determining step. This suggests

that a P(II) species may have a transient existence.

BIBLIOGRAPHY

BIBLIOGRAPHY

l. J. W. Mellor, Comprehensive Treatise on Inorganic and
Theoretical Chemistry, Wiley, (1956).

2. A. D. Mitchell, J. Chem Soc., 117, 1322(1920);
q12,0, 1266(1921). 1.31. 1624(192’2‘); 332' 629(1923).

3. R. O. Griffith and A. McKeown, Trans. Faraday Soc.,
43Q. 530(1934).

4. P. Haywood and D. M. Yost, J. Amer. Chem. Soc., 71,
915(1949). -

5. K. Pan and S.-H. Lin, J. Chinese Chem. Soc., Z, 75(1960).

6. S. K. Mishra and Y. K. Gupta, J. Inorg. Nucl. Chem.,
@2. 1643(1967).

7. K. S. Gupta and Y. K. Gupta, J. Chem. Soc. (A),
1970’ 256.
N'MA

8. W. A. Jenkins and D. M. Yost, J. Inorg. Nucl. Chem.,
11. 308(1959).

9. A. Fratiello and E. W. Anderson, J. Amer. Chem. Soc.,
23. 519(1963).

10. R. L. Carroll and L. B. Thomas, J. Amer. Chem. Soc.,
gg. 1376(1966).

11. J. N. C00per, H. L. Hort, C. W. Buffington, and
C. A. Holmes, J. Phys. Chem., 75, 891(1971).

12. J. N. Cooper, J. Phys. Chem., 34, 955(1970).

13. A. Sieverts and F. Loessner, Z. Anorg. Chem., 33,
1(1912).

14. G. C. R0per, T. E. Haas, and H. D. Gillman, Inorg. Chem.,
2. 1049(1970).

15. E. Ben-Zvi, Inorg. Chem.,.é, 1143(1967).

l6. E.66. Luoma, Ph.D. Thesis, Michigan State University,
19 .

17. K. Ogawa, Bull. Chem. Soc. Jap., 42, 1449(1969).

4O

18.

19.
20.
21.

22.

23.

24.

25.

26.

41
R. T. Jones and E. H. Swift, Anal. Chem., 35, 1272(1953).
R. P. Larsen and L. E. Ross, Anal. Chem., 31, 177(1959).
See Appendix 1

P. M. Beckwith and S. R. Crouch, Anal. Chem., 44,
221(1972). N

E. Stahl. Thin-Layer Chromatography, Academic Press,
New York, 1965, p. 482.

J. L. Dye and V. A. Nicely, J. Chem. Educ., 48,
443(1971)

A. G. Miroshnichenko and V. A. Lunenok-Burmakina,
Russ. J. lggrg. Chem, 15, 1345(1970).

P. R. Bevington, Data Reduction and Error Analysis
for the Physical Sciences, McGraw-Hill. New York,
1970.

R. K. Murmann, J. Amer. Chem. Soc., 92, 4184(1971).

APPENDIX A

APPENDIX A

Consider the rate equation

% == k2(Ru - x)( - g) (l)

where Ru and P represent initial concentrations of RuOu' and
H2P02-, respectively. The variable x corre3ponds to the
decrease in RuOu' concentration (or increase in RuOl,= con-
centration) due to the reaction with H2P02-. Data are Ruou'
concentration and time. Blank eXperiments (without H2P02-)
indicate reduction of RuOu' (attributed to contamination of
KOH pellets). Thus, when HZPOZI is added to RuOu- solution,
the production of RuOL,= is from two sources, and x in

equation (1) cannot be identified as the observed increase

in RuOL,= concentration. Clearly, x must be a value slightly
less than the observed increase in Ru0u= concentration at each
time t. In Figure AAL curve 1 is the case when RuOu' is not
reduced in a blank experiment and curve 2 represents reduction
of RuOu- in a blank. Curve 5 is the observed reduction of
RuOu' due to HZPOZ- and contamination. Curve 3 is curve 5
plus curve 2. The reduction of RuOu' due to HZPOZI is
represented by curve 4, which must lie between curves 3 and 5.
The variable x at any time t is the value of curve 4 minus the

value of curve 1. The problem is how to calculate curve 4.
i

ii

33.85.0331 .959 MM sogooom mo Psopxm

H-< mesmea

INITAI. [Ruo'i

[Ru 0;]

ARBITRARY
UNITS

iii

 

 

 

 

 

]
2

\ _
§>:;__3
5

 

 

 

TIME (ARBITRARY UNIT)

FIGURE
A-l

APPENDIX A (cont'd)

Ideally, the reduction of RuOu' due to the contamina-
tion would be treated as a concurrent reaction and the
individual rate expressions combined into a comprehensive
rate law. Without knowledge of the nature of the contamina-
tion or its concentration, such a mathematical eXpression
is difficult to justify (i.e., the stoichiometry).

It is reasonable to assume that, in curve 5, the
contribution due to contamination can be related to curve 2.
If, at time t, the blank eXperiment reveals that 5% of the
initial RuOu' has been reduced to Ru0u=, then 5% of the
observed decrease in RuOu' concentration in curve 5 is due
to the contamination. Thus, dividing the observed RuOu-
on curve 5 by 0.95 produces a new RuOu' concentration at
time t. Substracting this RuOu' concentration from curve 1
(the initial RuOu' concentration) gives the change in
Ruou' concentration due to the reaction with HZPOZI.

This value is x in equation (1).

The data were selected from eXperiments at 14.7. in
0.300 M KOH with Hzpoz' concentration of 2.49 x 10“ M
(Table V). In the first experiment 20.0 ml of RuOu'

solution was added to 20.0 ml of 0.300 M KOH (Table A-l).

iv

Time
(sec)

 

30
40

50

80

90
100
120
140
160
180
200
250
300
350
400
500
600

Time

 

120
250
tion

510
4110

Table A-1

Absorbance at 14.7°C, Experiment 1

 

Abgorbance [RuOqu
3 3.5 nm
1465.0 nm M x 10
0.502 1.724
0.506 1.718
0.522 1.697
(1.950) 1.667
(1.938) 1.649
0.550 1.659
0.554 1.653
(1.930) 1.637
0.563 1.641
0.566 1.637
0.568 1.634
0.570 1.632
(1.910) 1.606
0.598 1.594
0.600 1.591
(1.900) 1.591
0.630 1.551
Table A-2

[RuOQi/l .806

 

0.955
.951
.940
.923
.913
.919
.915
.906
.909
.906
.905
.904
.889
.883
.881
.881
.875
.859

TOTAL Ru Concentration, EXperiment 1

 

Absorbance
383.5 nm 465.0 nm .
1.929 0.560
1.910 0.580
1.000 0.600
1.893 0.615
1.890 0.621

1.639
1.610
l.<91
1.577
1.571

4

M,x 10

 

, RUOu-I

0.399
0.426

0.452
0.471

Q.h79

2

2.038
2.036
2.043

2.049
9.050

‘2 [043' 1

 

.006

APPENDIX A (cont'd)

Next, the eXperiment was repeated with 20.0 m1 of
2.49 x 10’“ M "7P02- in place of +he KOH solution (Table A3).
The first experiment was then repeated (Table A5). In
succeeding eXperiments the two solutions were alternated
in the described manner. The total RuOu' + Ru04= concen-
tration was calculated from absorbances in each eXperiment
by means of absorptivities given by Luomal6 (Tables A2, A4,
and A6).

Plots of RuOu' concentration Mg time made from data in
Tables A1 and A5 indicate reduction of RuOu- to Ru0u=.
This was always observed, never eliminated, and was
independent of care in cleaning vessels or care in prepara-
tion of reagents. The extent of the reduction decreased with
decreasing concentration of KOH as in eXperiments performed
to determine pH dependence. It is reasonable to conclude the
KOH pellets contained a trace of oxidizable matter.

The initial [Ruou'] (at mixing) is obtained by
extrapolation of the smoothed curve of [huO4I] Mg time from

blank eXperiments. (A computer program written for inter-

polation and extrapolation25 was used for this purpose.

Results are listed in Table A7.) This extrapolated value

was the basis for calculating the fraction of [Ruou-J
vi

Table A93

Absorbance at 14.7°C, Experiment 2

 

 

 

 

 

 

 

M x 10“
Time) Abgorbance [11qu [Ruou'] g = _
sec 3 3.5 nm 1. 07- [1in ]
465.0 nm Factor 4
30 0.890 1.243 1.293 0.514
40 0.990 1.108 1.159 0.648
60 1.158 0.881 0.947 0.860
70 1.223 0.793 0.854 0.953
80 1.282 0.713 0.775 1.032
90 1.320 0.662 0.720 1.087
100 1.365 0.601 0.657 1.150
120 (1.186) 0.477 0.526 1.181
140 1.514 0.399 0.441 1.366
160 1.563 0.333 0.369 1.438
180 1.602 0.281 0.312 1.495
200 1.637 0.233 0.259 1.548
250 (0.967) 0.143 0.162 1.645
300 (0.932) 0.089 0.101 1.706
50 1.769 0.0 5 0.063 1.794
00 1.780 0.0 0 0.046 1.761
500 1.805 0.006 0.007 1.800
600 1.816 0 0 1.807
Table .A-4
TOTAL Ru Concentration, EXperiment l
M x 10“
Absorbance -
Time 383.5 nm 465.0 nm lRuOu'I RuOu- Z
140 1.139 1.515 0.403 1.681 2.084
230 0.981 1.680 0.168 1.906 2.074
320 0.921 1.756 0.072 2.008 2.080
480 0.882 1.802 0.012 2.070 2.082
550 0.864 1.810 0.014 2.092 2.078

vii

 

2.080 1 .006

Table A-5

Absorbance at 14.7°C, EXperiment 3

 

 

 

 

Time Absorbance [Ruou'l [111104.] Mean col 4
(sec) (383.5 mm) M x 10 1 809 Tables 11
465.0 nm ' ___&5___and
0 0.512 1.751 0.968 '0.961
0 0.522 1.738 0.961 .956
50 0.545 1.707 .944 .942
60 (1.985) 1.698 .939 .930
70 (1.980) 1.690 .934 .929
80 (1.972) 1.678 .928 .920
90 (1.964) 1.666 .921 .920
100 (1.955) 1.652 .913 .914
120 0.594 1.640 .907 .907
140 0.600 1.632 .902 .905
160 (1.940) 1.629 .901 .903
180 (1.936) 1.623 .897 .901
200 (1.933) 1.619 .895 .899
250 0.630 1.592 .880 .885
300 (1.915) 1.591 .880 .881
350 (1.912) 1.587 .877 .879
400 0.646 1.570 .868 .874
500 (1.900) 1.568 .867 .871
600 (1.893) 1.558 .861 .845
Table A-6
TOTAL Ru Concentration, EXperiment 3 4
Absorbance M_x 10
.__________ _ ,___1:__
Time 383.5 nm 465.0 nm lRuOu I 111110,+ Z
160 1.940 0.604 1. 2 0.451 2.079
290 1.915 0.631 1.591 0.488 2.079
520 1.879 0.653 1.565 0.503 2.068
600 1.893 0.662 1.554 0.529 2. 083
750 1.885 0.671 1.542 0.541 2.083

 

2.078 1 .006

viii

Table .A-7

Extrapolated Ruou- Concentration at t = 0.0 sec.

 

 

 

Data from
Table 0 _30 40 50
Al 1.806* 1.731 1.710 1.691
A1 1.809 1.752 1.733 1.714

[Ruou'] M x 10

ix

APPENDIX A (cont'd)
remaining at each time t (column 4 of Tables A1 and A5
for eXperiments preceeding and following the H2P02' reaction.
These values were averaged (Table A5, column 5) and divided
into the observed [RuO4-] for the HZPOZ- reaction (Table A3,
column 4). The input variable x is obtained by subtracting
the corrected Eiuou'] (Table A3. column4) from the average
of the extrapolated initial [éuou'] (Table A7). Figure A2
is a computer plot of x (input and calculated from equa-
tion (2)) Kg time for data in Table A3 with initial
{Ruouj = 1.807 x 10'“ m. If the initial [Ruou'] is
allowed to be an adjustable parameter, the calculated
value is usually within 3% of the extrapolated value
(Figure A3). In this case, the values of Em04-J and k2

L;

which minimize the residuals are 1.829 x 10' M_and

47.3 mil sec-1. respectively, as Opposed to 1.807 x 10'“ M

and 48.5 M81 sec-l.

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APPENDIX B
OTHER EXPERIMENTAL WORK

APPENDIX B

OTHER EXPERIMENTAL WORK

The reduction of Tl(III) by hypoPhosphorus acid in
perchloric acid was the first project undertaken. All
reagents were purified and stock solutions prepared.
Before experiments were begun, the published results of

Gupta and Gupta7

were received. The project was abandoned
and efforts were redirected along other lines. It was
noted that white crystals were obtained from concentrated
solutions of Tl(III) and H3P02. These crystals decomposed
to a yellow colored substance with a phosphine smell when
they were dried.

The oxygen necessary to convert P(I) to P(III)
usually originates in the solvent water. If the oxident
also has oxygen, from which source would the oxygen come?
The second project was directed to this question. The
search for a reducible (by hypOphOSphite at a measurable
rate) oxocation included the following Species:

URANYL ION

No change was observed in the proton nmr of H3P02

after three days in a solution 0.2 M in H3P02 and U02+.

XV

APPENDIX B (cont'd)
The solution was brought to boiling and C12 gas (from
reduction of C104“) was liberated. Substituting N03' for

010“- and heating produced lower oxides of nitrogen.

Finally, peroxouranyl (U04) was dissolved in HCL solution.

and heated to form 0022*. After adding HBPOZ, the solution
was purged of oxygen and allowed to reflux overnight. Again,
no change in H31302 concentration was observed in the proton
nmr. Yellow plates precipitated from concentrated (about

2+ and H

1 M) solutions of U02 3

P02.
VANADYL ION
A color change (blue to green) was observed when

2+ and H3P02 were mixed. COOper12 reported

solutions of V0
this to be a complex formation rather than a redox reaction.
PEROXOTITANYL

This yellow colored oxocation (Ti022+) was not noticably
reduced in 24 hours at room temperature. The study was not
thorough and should be reinvestigated.
SODIUM PERRUTHENATE

A substantial amount of data.were collected with the sodium

salts of hypophosphite and perruthenate. These data were

seldom reproducible. It is likely the problem originated
xvi

APPENDIX B (cont'd)
in the preparation of NaRuOu. When RuOu is generated

from NaRiO and RuClq, 01' is also oxidized to 01?.

3
When this dissolves in alkali hypochlorite, 001-, is

produced. This powerful oxident oxidizes RuOu= to RuOu-
and RuOu. (In at least one experiment RuOu' was reduced
by H2P02- to RuOu= which was oxidized back to RuOu', all

within 20 minutes.) Several attempts to crystallize

NaRuOu were unsuccessful.

xvii