REDUCTION OF SILVER AMINE COMPLEXES
BY CARBON MONOXIDE
by
SHUZO NAKAMURA
Bo Sc. i n Engineering, Kyoto University, 1960
A THESIS SUBMITTED IN PARTIAL FULFILMENT OF
THE REQUIREMENTS FOR THE DEGREE OF
MASTER OF SCIENCE
i n the Department of
CHEMISTRY
We accept t h i s thesis as conforming to the
required standard
THE UNIVERSITY OF BRITISH COLUMBIA
August, 1962
In presenting this thesis in partial fulfilment of
the requirements for an advanced degree at the University of
British Columbia, I agree that the Library shall make i t freely
available for reference and study.
I further agree that permission
for extensive copying of this thesis for scholarly purposes may be
granted by the Head of my Department or by his representatives.
It is understood that copying or publication of this thesis for
financial gain shall not be allowed without my written permission.
Department
The University of British Colombia,
Vancouver 8, Canada.
ii
ABSTRACT
The k i n e t i c s of the reduction of s i l v e r amine complexes
i n aqueous solution by carbon monoxide were investigated.
For a number of amines including e t h y l - , methyl-, d i e t h y l ethanol-, diethanolamine and some primary diamines, the
rate law was found to be of the form;
d(Ag(I))
"
dt
[AgL 3[CO]
+
d(CO)
~ ~~dt~
=
2
=
2
k e x
P
[LH+j
- 2k (L-Ag-OH ]tCOj
where L denotes the amine.
(1)
These k i n e t i c s were interpreted
i n terms of the following mechanism.
AgL
+
2
+ H 0 ^ r = ± : L-Ag-OH + LH+ (Rapid)
2
(2)
k
L-Ag-OH + CO
L-Ag-COOH (Rate-determining)
L-Ag-COOH + Ag(I)
>• Products (Rapid)
(3)
(4)
The rate constant of the rate-determining step (3) was found to
be nearly independent of the nature of the amine molecule, L,
coordinated to s i l v e r ion, using the b a s i c i t y constants of the
amines and d i s s o c i a t i o n constants of the corresponding s i l v e r
amine complexes. The actual o v e r a l l rate of the reaction
varied with the nature of amine but t h i s was a t t r i b u t a b l e only
to the d i f f e r e n t equilibrium concentrations of L-Ag-OH. The
rate of t h i s rate-determining bimolecular process was found
to be s u r p r i s i n g l y fast; k__o = 5x10 mole
1. sec. ,
iii
AH
^
9 K c a l . mole"
s i l v e r i o n by CO
and
AS
—
-15
e,u.
The r e d u c t i o n o f
i n a c i d i c o r n e u t r a l media i s known to be
slow and t h i s can now
very
be a t t r i b u t e d to the base c a t a l y z e d
nature o f the r e a c t i o n .
S i l v e r complexes o f primary diamines
(ethylenediamine,
1,3-diaminopropane, e t c . ) were reduced more s l o w l y ; t h i s
was
a t t r i b u t e d to the s t a b i l i z a t i o n o f mono-complexed s i l v e r
(I)
s p e c i e s by c h e l a t e f o r m a t i o n .
I n the case o f ammonia normal k i n e t i c s were observed
at
h i g h e r pH but a t Lower pH the r a t e became second o r d e r i n
(Ag(I))and i n v e r s e l y second o r d e r i n (NH^ ).
T h i s was
+
buted to c o m p e t i t i o n between decomposition
o f the i n t e r m e d i a t e
complex and i t s f u r t h e r r e a c t i o n w i t h another Ag(I)
to g i v e m e t a l l i c s i l v e r and carbon d i o x i d e .
s i m i l a r competition was
found w i t h two
t r i e t h y l a m i n e and t r i e t h a n o l a m i n e .
attri-
species
Evidence f o r
t e r t i a r y amines, i . e . ,
ACKNOWLEDGEMENTS
The author wishes t o express h i s s i n c e r e g r a t i t u t e f o r
the c o n t i n u i n g a d v i c e , h e l p and encouragement g i v e n by Dr. J .
Halpern, who suggested and d i r e c t e d t h i s study, and t o Dr. E.
P e t e r s and Mr. R. T. McAndrew o f t h e Department o f M i n i n g and
M e t a l l u r g y f o r i n f o r m a t i o n about t h e i r r e l a t e d work.
He i s a l s o g r a t e f u l t o Dr. C. A. McDowell, Head o f t h e
Department o f Chemistry, who enabled him t o work i n t h i s
Department.
Support o f t h i s work by the A l f r e d P. S l o a n Foundation
and the N a t i o n a l Research C o u n c i l o f Canada i s a l s o
acknowledged.
gratefully
V
TABLE OF CONTENTS
Page
I.
II.
INTRODUCTION
1
EXPERIMENTAL. .
6
MATERIAL.
6
ANALYSIS
7
PROCEDURE
K i n e t i c Measurements
Stoichiometry Measurements
III.
RESULTS AND DISCUSSIONS
.7
7
11
13
STOICHIOMETRY
13
KINETICS AND MECHANISM.
18
Ethylamine Complex.
19
Other "Standard" Systems
25
Triethylamine Complex
33
Triethanolamine Complex
40
Ammonia Complex . . . . . . . .
47
Diamine Complexes
56
GENERAL DISCUSSION
REFERENCES
APPENDIX I .
64
71
Selected Thermodynamic Properties of Amines
and S i l v e r Amine Complexes
72
vi
LIST OF TABLES
Table No.
Page
lo
Results of Stoichiometry Measurements (I)
2.
Results of Stoichiometry Measurements (II)
. 1 4
3o
Results of Stoichiometry Measurements (III)
. 1 8
4.
Rate of Reaction of Ethylamine and Related Amine
Complexes of S i l v e r
5.
1 3
.
.
.
.
.
2 2
Summary of K i n e t i c and Related Thermodynamic Data f o r
"Standard" Systems at 25°C. .
6.
.
. . . . . 3 0
Apparent Enthalpy and Entropy of A c t i v a t i o n f o r
"Standard" Systems
3 2
7.
Rate of Reaction of Triethylamine Complex at
8.
Rate of Reaction of Triethanolamine Complex . . . . . . 4 1
9.
Rate of Reaction of Ammonia Complex . . . . . . . . . . 4 8
10.
Rate of Reaction of Diamine Complexes . . . . . . . . . 57
11.
Rates of Diamine Complexes and Their S t a b i l i t y
12.
Summary of Kinetic and Related Thermodynamic Data . . . 65
25°C.
.
.
3 4
vii
LIST OF FIGURES
Figure No.
I.
II.
Page
Gas bubbling glass apparatus.
8
T y p i c a l t i t r a t i o n curves f o r f i n a l reaction
solutions . . . . .
III.
IV.
16
Typical rate p l o t s f o r ethylamine complex
Dependence of rate on carbon monoxide
concentration
at 25°C. f o r ethylamine complex
V.
20
21
Dependence of rate on ammonium ion concentration at
25 C. f o r ethylamine complex
24
VI.
Typical rate plots f o r ethylamine-type complexes. . . 26
VII.
Arrhenius p l o t s f o r ethylamine-type complexes . . . . 31
VIII.
T y p i c a l rate p l o t s f o r triethylamine complex. . . . . 35
IX.
Dependence of rate on free amine concentration at
25°C. f o r triethylamine complex .
X.
XI.
. 38
T y p i c a l rate plots f o r triethanolamine
Arrhenius
plots f o r triethanolamine
complex. . . . 42
and ethylene-
diamine complexes
XII.
X I I I .
46
T y p i c a l rate plots f o r ammonia complex.
Dependence of rate on ammonium ion concentration at
30°C, f o r ammonia complex
XIV.
XV.
XVI.
50
51
Arrhenius plot f o r ammonia system
54
Typical rate plots f o r diamine complexes
58
Dependence of the rate on free amine
at 25°C. f o r 1,3-diaminopropane
concentration
60
I.
INTRODUCTION
Recently, molecular hydrogen
s
although unreactive toward
the majority of common inorganic o x i d i z i n g agents, was
found
to be oxidized under r e l a t i v e l y mild conditions i n aqueous
2+
solution by a few metal ions and complexes, notably Cu
Ag ,
Hg2 , Hg^s,
+
and MnQ^~„
+
s
Halpern and h i s coworkers (1)
have studied these systems extensively and have elucidated
the mechanism through which the r e l a t i v e l y strong H-H
-1
having a d i s s o c i a t i o n energy of 103 Kcal. mole
bond,
, i s activated,
and the dependence of the r e a c t i v i t y on the electron configurat i o n of the central metal ions.
These studies have prompted
s i m i l a r studies on other i n e r t reducing agents.
The mechanism
of these reactions and the nature of metal ions and complexes
which a c t i v a t e those inert molecules are of great i n t e r e s t
for the study of chemical r e a c t i v i t y , i n general, and e s p e c i a l l y
of the c a t a l y t i c a c t i v i t y of t r a n s i t i o n metals and t h e i r compounds.
Amongst other r e l a t i v e l y inert reducing agents whose reactions have been investigated i n t h i s laboratory are carbon
monoxide and formic a c i d .
-
MnO^
s
94-
Hg
, Hgj
2+
The reactions of the l a t t e r with
%4-
and T l
were studied by Taylor and Halpern
(2) and t h e i r k i n e t i c s and mechanisms were elucidated.
Harkness
and Halpern (3) examined the reaction of CO with those metal
ions which are active toward molecular hydrogen and found that
2
only Hg
and MnO^
showed measurable r e a c t i v i t y toward CO i n
homogeneous aqueous solution under moderate condition.
found F e ^
+
Tl"**" and Cr^O^ also to be i n a c t i v e .
-
9
At
They
elevated
temperature and pressure Bauch et a l . (4) observed that s i l v e r
sulfate and cupric s u l f a t e i n aqueous solution also were
reduced by CO..
reaction was
They reported that the rate of the former
second order i n Ag(I)
and was
the solution with ammonium acetate.
enhanced by buffering
However, they did not
study the dependence of the rate on pH.
Following
t h i s work,
and a f t e r commencement of the present study, Peters and McAndrew
(5)
studied the reaction of s i l v e r acetate i n aqueous solution
with CO i n further d e t a i l under experimental conditions s i m i l a r
to those of Bauch et a l . (4).
These r e l a t e d studies
are
summarized below.
2+
Hg
(3)...This i s the only metal ion which was
found
to oxidize CO i n aqueous solution under r e l a t i v e l y mild
condi-
tions (atmospheric pressure and below 8 0 C ) i n the absence
24of complexing agents. For the reduction of Hg
, i.e.,
2 H g + CO + H 0
HgJ + C0 + 2H+
(1-D
o
2+
+
2
2
k i n e t i c measurements i n d i l u t e HC10.
solutution over the
temperature range 26 to 54 C. yielded the pH-independent
rate
law
(1-2)
with AH
*
= 14.6
Kcal. mole
-1
and
AS
"k
= -13 e.u.
This
was
interpreted i n terms of the following mechanism.
4-
0
-Hg
z+
0H + CO
*- -Hg-C-OH
2
-Hg-C-OH
+ H
> Hg + C0 + H
2
Hg + H g
Hg
2 +
2 +
(slow)
(1-3)
(fast)
(1-4)
(fast)
(1-5)
Support for the proposed intermediate complex i s provided by the
P,
i s o l a t i o n of a stable analogue, AcO-Hg-C-OCH^, formed by reaction
of CO with mercuric acetate i n methanol solution (6).
Mn0 ~
(3)...The reduction of MnO^" by CO (to Mn0 i n
4
2
a c i d i c and neutral solutions and to MnO^
i n basic
solutions)
was found to proceed r e a d i l y over the temperature range 28 to
50°C.
The complete rate law was found to be
" ^dT
with
AH
=
k
CC0J[Mn0 ")
4
= 13 Kcal. mole" and AS
1
(1-6)
= -17 e.u., both s u b s t a n t i a l l y
constant over the pH range L to 13, which confirmed and extended
the e a r l i e r k i n e t i c measurements on t h i s system by Just and
Kauko (7). Harkness and Halpern (3) also found that t h i s system
shows a remarkable c a t a l y t i c e f f e c t on the addition of A g and
+
Hg
2 +
(but not C u , Fe *, C d , or T l
2 +
3
2 +
3 +
) which they a t t r i b u t e d
to favorable reaction paths involving intermediate such as
Ag-CO-OMnOg
Hg
2 +
with CO.
analogous to that postulated i n the reaction of
4
Ag SO^ and CuSO^
2
(4)...Bauch et a l . studied the aqueous
Ag SO^ system over the temperature range 70 to 110°C. under
2
CO pressure up to 50 atmosphere and found the reaction
Ag
+
+ %C0 + %H 0
> Ag + %C0
2
+ H
2
(1-7)
+
proceeded according to the rate law given by (1-8).
" ^df^
=
^SxlO
Ug )
+
6
2
P
c o
e-
1 4
»
0 0
°/
R T
(1-8)
They also observed that the rate was increased by buffering the
solution with ammonium acetate, the k i n e t i c s f o r the buffered
s i l v e r sulfate system being given by (1-9).
-
= 6.02x10*
Ug )
+
2
P
e- '
9
C Q
3 0 0 / R T
(1-9)
They attributed this difference i n rate to the favorable
dependence of the equilibrium on increasing pH, and proposed
the same mechanism for both buffered and unbuffered system, i . e . ,
Ag
+
+ CO 5 = = ^ Ag(C0)
Ag(C0) + A g
+
Ag (C0)
4 +
2
+ H0
2
+
+
(rapid e q u i l i b r i a ) (1-10)
~ = r Ag (C0)" "
H
2
>2Ag + C0
2
+ 2H
+
(rate determining) (1-11)
In support of t h i s mechanism, they c i t e d the existence of
carbonyl complex of the " f i r s t subgroup" such as (Cu(Cl,Br)C0)^,
Ag (C0)S0^ and (AuCl'CO)^.
2
I t i s obvious from t h e i r r e s u l t s
that there must be some pH-dependent process contributing to the
o v e r a l l rate but t h i s was not elucidated.
For the reduction of
CuSO^ by CO they also observed second order dependence of the
rate on [Cu
), the rate over the temperature range 160 to 190 C.
being given by
3
d ICu
dt
2.56X10
13
(Cu^)
2
*
e
"
3 3
'
5 0
°/
(1-12)
R T
In t h i s system the e f f e c t of b u f f e r i n g was not reported because
of experimental
difficulties.
Recently, Peters and McAndrew (5) have extended t h i s work
on the reduction of s i l v e r s a l t s i n a c i d i c solution.
Both i n
acetate-buffered and perchlorate media the rate was found to
be very slow, requiring the use of elevated temperature (>90°C.)
and CO pressure (10 to 30 atm.).
The r e s u l t s of t h i s work w i l l
be considered l a t e r .
The present study i s concerned with the reduction of s i l v e r
amine complexes by CO i n basic media.
In contrast to the behaviour
i n a c i d solutions the reaction under these conditions i s rapid
and r e a d i l y measureable at room temperature and atmospheric
pressure.
6
II.
EXPERIMENTAL
MATERIALS
S i l v e r perchlorate was G. F. Smith Reagent grade and was
unaffected by r e c r y s t a l l i z a t i o n .
and Adamson 60% Reagent grade.
Perchloric a c i d was
Baker
Ethylenediamine, Fisher c e r t i -
f i e d reagent, was used without further p u r i f i c a t i o n .
Distilla-
t i o n of t h i s product had no e f f e c t on the reaction rate.
Matheson triethylamine, which contained a reducing impurity,
was p u r i f i e d by passing through a molecular sieve column and
then d i s t i l l e d under 120 mm.
Hg nitrogen atmosphere.
Matheson
33% aqueous solution of ethylamine, and diethylamine (b.p.
55-56°C); B.D.H. 25/30% methylamine aqueous solution, pureethanolamine, diethanolamine and triethanolamine, and K & K
Laboratories' 1,3-diaminopropane were used without further
purification.
K & K Laboratories' 1,4-diaminobutane was
t i l l e d at 20 mm.
Hg before use.
redis-
Ordinary d i s t i l l e d water was
used i n the preparation of a l l solutions and gave rates i d e n t i c a l
with those obtained with water d i s t i l l e d from a l k a l i n e permanganate.
Nitrogen gas was supplied by the Canadian L i q u i d A i r Co.
Carbon monoxide ( C P . grade) and CO-^
from Matheson of Canada L t d .
gas mixtures were obtained
The chromatographic analysis of
a l l these gases revealed substantially no contamination by
oxygen.
In a l l experiments the amine perchlorate was prepared
7
by n e u t r a l i z i n g the amine w i t h p e r c h l o r i c a c i d .
The e x p e r i m e n t a l
s o l u t i o n s were prepared by d i l u t i n g a l i q u o t s o f s t a n d a r d i z e d
stock solutions.
ANALYSIS
The n o r m a l i t y o f amines and aqueous amine s o l u t i o n s
was determined by t i t r a t i o n w i t h standard h y d r o c h l o r i c a c i d .
S i l v e r i o n c o n c e n t r a t i o n was determined by t h i o c y a n a t e t i t r a t i o n
in acidic solution with f e r r i c indicator.
Carbon monoxide and
n i t r o g e n gas mixtures were a n a l y z e d w i t h a Beckman GC-2 gas
chromatograph
u s i n g a m o l e c u l a r s i e v e column.
PROCEDURES
K i n e t i c measurements
Except f o r the ammonia system, r a t e s o f
a l l the r e a c t i o n s were determined a t atmospheric p r e s s u r e , by
b u b b l i n g the CO gas ( o r a CO-N^ mixture) through the s o l u t i o n
i n the g l a s s apparatus d e p i c t e d i n F i g u r e I .
The gas was passed
through a p r e s a t u r a t o r f i l l e d w i t h aqueous s o l u t i o n o f NaNOg
and the amine t o e s t a b l i s h the same p a r t i a l p r e s s u r e o f water
and the amine as the r e a c t i o n s o l u t i o n , and was then d i s p e r s e d
through a s i n t e r e d g l a s s p l a t e i n t o the r e a c t i o n
solution.
The e f f l u e n t gas was l e d t o a gas flame and was burned.
whole apparatus was immersed i n a c o n s t a n t temperature
thermostated t o ~ 0.03°C.
The
bath
I t was e s t a b l i s h e d t h a t the f l o w r a t e
o f the gas d i d n o t a f f e c t the observed r e a c t i o n r a t e ; hence i t
may be assumed t h a t the s o l u t i o n s were s a t u r a t e d w i t h the gas.
Fig. I.
Gas Bubbling
Glass Apparatus
A.
Gas I n l e t
B.
Presaturater
C.
Sintered Glass P l a t e
D.
Reaction
E.
Gas Outlet
F.
Gas O u t l e t Stopper
G.
Sampling Tube
H.
Sample Outlet
Solution
Mixture
CO
9
A f t e r placing 250-500 ml. of the reaction mixture of the
desired composition i n the apparatus, the system was allowed
to a t t a i n thermal equilibrium under nitrogen flow.
The s t a b i l i t y
of the reaction mixture was checked by sampling and analyzing
the solution several times under the nitrogen flow and then
the gas flow was switched to CO or to a CO-^ mixture.
The
solution was sampled p e r i o d i c a l l y and the samples were analyzed
as described previously.
The time required f o r saturation of
the solution with the gas was usually n e g l i g i b l e ; less than
30-60 sec. For reaction solutions i n which the t o t a l solute
concentrations were lower than 0.5-0.6 molar, the p a r t i a l
pressure of the gas was assumed to be atmospheric pressure
minus the vapor pressure of pure water at the reaction temperature.
In the case of triethylamine, which required a very high
amine concentration to obtain stable solutions, Lattey°s (8)
data f o r the t o t a l vapor pressure of triethylamine-water mixtures
were used.
Variation of the CO p a r t i a l pressure was achieved,
when desired, by using analyzed C O - N 2 mixtures.
With ammonia, whose p a r t i a l pressure i s very high and
whose reaction rate was very low at atmospheric pressure,
an autoclave was used.
The apparatus used was a Parr Series 4500
autoclave with a g l a s s - l i n e d stainless s t e e l reaction vessel,
provided with a s t i r r e r , gas i n l e t tube, sampling tube f i t t e d
with a stainless s t e e l f i l t e r , pressure gauge and thermowell,
surrounded by an e l e c t r i c heating mantle controlled by a rheostat.
10
F i n e temperature c o n t r o l was
a c h i e v e d by use o f an a u x i l i a r y
e l e c t r i c h e a t e r , immersed i n the r e a c t i o n m i x t u r e through the
g l a s s - l i n e d thermowell, and c o n t r o l l e d by a Thermistemp Temperat u r e C o n t r o l l e r (Model 71) a c t u a t e d by a t h e r m i s t o r immersed
i n the s o l u t i o n .
T h i s arrangement
gave temperature
control
of d= 0.3°C,
A 1 500 ml.
s
r e a c t i o n m i x t u r e was made up from s t o c k
s o l u t i o n s and p l a c e d i n the r e a c t i o n v e s s e l .
was
N i t r o g e n gas
run i n t o the m i x t u r e through the sampling tube and the
porous s t a i n l e s s s t e e l f i l t e r , under a g i t a t i o n by the s t i r r e r ,
f o r some f i v e minutes; then the v e s s e l was
to the d e s i r e d temperature.
was
s e a l e d and brought
The s t a b i l i t y o f the s o l u t i o n
e s t a b l i s h e d by t a k i n g samples and a n a l y z i n g them f o r s i l v e r
i o n over a one hour p e r i o d .
reduced to one atmosphere
the CO gas was
The i n t e r n a l p r e s s u r e was
then
by opening the gas o u t l e t once,
and
i n t r o d u c e d from a CO c y l i n d e r and m a i n t a i n e d a t
a desired pressure.
The p a r t i a l p r e s s u r e o f CO was
as the d i f f e r e n c e between the t o t a l p r e s s u r e and the
calculated
combined
vapor p r e s s u r e o f the s o l u t i o n and r e s i d u a l n i t r o g e n p r e s s u r e .
Samples were taken a t a p p r o p r i a t e i n t e r v a l s and a n a l y z e d f o r
silver ion.
A f t e r each sampling more CO gas was
supplied to
* The nichrome w i r e c o i l o f the a u x i l i a r y h e a t e r was made
to occupy the lower h a l f o f the thermowell and a volume o f
s o l u t i o n s u f f i c i e n t t o f i l l the g l a s s l i n e r was used so as
to a v o i d u n d e s i r a b l e s u p e r h e a t i n g o f the thermowell.
11
the system to e s t a b l i s h a constant pressure.
The s t i r r e r
was
rotated, usually at 600 r.p.m., a f t e r establishing that the
rate of reaction was
Stoichiometry
independent of the s t i r r i n g rate.
measurements
The amount of CO gas consumed was
measured by means of a gas burette apparatus.
conical
f l a s k reaction vessel was
A 125
ml.
placed i n a small poly-
ethylene water bath (diameter ca. 5 in.) which was
connected
to a thermostated water c i r c u l a t o r . The reaction v e s s e l f l a s k
was
connected to a 50 c c . burette provided with a mercury
balancing b o t t l e .
The reaction solution was
teflon-coated magnetic s t i r r e r .
was
s t i r r e d with a
100 ml. of reaction solution
placed i n the reaction vessel; nitrogen gas was
first
passed through the solution to remove oxygen from the reaction
system, and f i n a l l y a f t e r d i s p l a c i n g the nitrogen by CO
f i l l i n g the reaction vessel with CO,
the whole system
closed, the gas burette mercury l e v e l was
magnetic s t i r r e r was
turned on.
balanced and
and
was
the
A dibutylphthalate a u x i l i a r y
manometer was used for f i n e adjustment of the mercury l e v e l .
When the rate of gas uptake became s u f f i c i e n t l y slow the
burette was
read, the reaction vessel was
the burette and the solution was
gas uptake was
and
analyzed.
disconnected from
The volume of
corrected f o r the s o l u b i l i t y of CO,
temperature
pressure.
In several cases the reaction mixture was
t i t r a t e d for
12
amine and carbonate before and a f t e r the reaction.
In those
cases, an o r i g i n a l reaction solution was flushed with carbon
dioxide-free nitrogen gas before i t was placed into the reaction
vessel.
A 25 ml. aliquot was taken for analysis and 100 ml. was
pipetted under
into the reaction v e s s e l .
The o r i g i n a l and
the f i n a l solutions were analyzed for amine and carbonate by
potentiometric t i t r a t i o n with standard hydrochloric acid using
a Beckman pH meter.
S i l v e r metal deposited i n the reaction vessel was
on a glass c r u c i b l e and weighed.
collected
I t s purity was determined by
d i s s o l v i n g i n n i t r i c a c i d and t i t r a t i n g with thiocyanate.
The amount of carbonate i n the f i n a l reaction mixture
was determined more d i r e c t l y by a gravimetric method, as
barium carbonate.
Barium n i t r a t e or barium hydroxide was
used to p r e c i p i t a t e the carbonate.
barium hydroxide was used.
For weakly basic amines
In some cases ethylenediamine or
propylenediamine was added to s t a b i l i z e the remaining s i l v e r
ion before barium hydroxide was added.
13
III.
RESULTS AND DISCUSSIONS
STOICHIOMETRY
For a l l the amines which were used i n the present study,
when aqueous s i l v e r (I) amine solution was reacted with CO,
a brown or black s o l i d separated from the solution, which
proved on analysis to be pure s i l v e r metal.
In many cases
the glass w a l l of the reaction vessel also was covered by
a s i l v e r mirror.
The amount of s i l v e r metal formed was found
to be equal to the amount of s i l v e r ion reduced (Table 1).
TABLE 1
RESULTS OF STOICHIOMETRY MEASUREMENTS (I)
Amine
Initial
(Ag(I)]
mole/1
Amine
PerSilver
chlorate Amine Reduced
mole/1 mole/1 mole/1
Carbon- S i l v e r
Monoxide Metal
Uptake Recovered
mole/1 mole/1
0.0299
0.000
0.300
0.0299
0.015
0.0295
1,3-diaminopropane
0.0297
0.000
0.300
0.0297
0.015
0.0295
dlethanolamine
0.000
0.300
0.0296
0.0129
0.0294
ethylamine
0.0296
14
TABLE 2
RESULTS OF STOICHIOMETRY MEASUREMENTS ( I I )
•1 •
2
I n i t i a l Amine
Carbon- Hydrogen Carbon(Ag(I))
PerS i l v e r Monoxide I o n
ate
c h l o r a t e Amine Reduced Uptake Produced Produced
mole/1 mole/1 mole/1 mole/1
mole/1 mole/1
mole/1
2
Amine
ethylamine
0.0300
0.000
0.0963
0.0300
0.015
0.0606 0.0151
methylamine
0.0299
0.000
0.0937
0.0299
0.015
0.0605 0.0150
diethylamine
0.0296
0.100
0.0933
0.0292
0.0146
0.0620 0.0145
ethanolamine
0.0300
0.000
0.100
0.0178
0.00809 ***
dlethanolamine
0.0296
0.000
0.0988
0.0250
3
3
0.0114 (0.0532)(0.0133)
0.0297
0.000
0.300
0.0296
0.0129 ***"
3
***
3
u 3
triethanolamine
0.0298
0.000 0.300
0.0199 0.0091
ethylenediamine
0.000
0.0298
0.0150
0.0618 0.0148
0.0297
0.015
***
0.0300
1,3-diaminepropane
0.0297
0.0648
0.000 0.300
***-
***
1 I n those experiments where t h e r e a c t i o n was c o n t i n u e d t o complet i o n a slow f u r t h e r uptake o f CO was observed even when a l l the
s i l v e r had been reduced. T h i s zero o r d e r uptake o f CO i s presumably
a t t r i b u t a b l e t o r e a c t i o n o f CO and H2O on t h e s i l v e r metal s u r f a c e
to form formate (9) «
>
Jj
2
These a r e t h e r e s u l t s from p H - t i t r a t i o n o f t h e f i n a l
solution.
3 I n t h e case o f ethanolamines, which a r e the l e a s t b a s i c o f a l l
these amines, pH t i t r a t i o n was h o t r e a d i l y a p p l i c a b l e . F o r
diethanolamine, t h i s method y i e l d e d o n l y the combined c o n c e n t r a t i o n o f amine and carbonate. These r e s u l t s i n t h e parentheses
were c a l c u l a t e d from t h i s assuming hydrogen i o n produced: /
carbonate = 4s1, hence s u b j e c t t o e r r o r .
'
15
Results of CO uptake measurements i n Table 2 show that
two gram-ions of s i l v e r ion were reduced f o r each mole of
*
CO .
The pH t i t r a t i o n of the f i n a l solution (results are
summarized i n Table 2 and t y p i c a l t i t r a t i o n curves are given
i n Figure II) showed that two gram-ions of hydrogen ion and
one-half gram ion of carbonate (identity of t h i s product i s
to be discussed l a t e r ) were produced f o r each gram-ion of s i l v e r
ion reduced by one-half mole of CO.
The o v e r a l l reaction can
therefore be represented by
2AgL
+
2
+ CO + 2H 0
2
s*2Ag + C0 ~ + 4LH
+
3
(3-1)
under conditions where the amine (represented by L) i s present
i n excess.
The r e s u l t s of d i r e c t carbonate determinations (gravimetric a l l y as barium carbonate) on the same f i n a l reaction solutions
are
shown i n Table 3.
ammonia systems
gave
the t i t r a t i o n r e s u l t s .
Only triethylamine, diethylamine and
those y i e l d s of carbonate expected from
In the other cases a portion of the
* In the case of the three ethanolamines the CO uptake was
about 10% lower than the t h e o r e t i c a l value (Table 2). In these
cases, some s i l v e r apparently a l s o was reduced by the amines or
by an impurity. These side reactions were most pronounced at
the high pH of these CO uptake experiments i n which; i n order
to increase the rate of reaction, no amine perchlorate was
added. Such solutions deposited some m e t a l l i c s i l v e r on
standing even i n the absence of CO, while solutions containing
amine perchlorate ( i . e . those used f o r the k i n e t i c experiments)
were stable. The amount of s i l v e r reduced by these side
reactions seemed to be d i r e c t l y dependent on the amine concent r a t i o n as indicated by the two experiments with diethanolamine i n Table 2.
5
io
15
20
25
30
0.1 U Standard HCl, ml.
Pig. I I . Typical T i t r a t i o n Curves of Pinal Reaction Mixtures
—'•—
See Table 2 f o r experimental conditions.
17
"carbonate", r e s u l t i n g from the oxidation of the CO apparently
combines with the amine to form a substance which decomposes on
acidification.
Decomposition with release of carbonate also
occurred on treatment with base.
Thus, when the f i n a l reaction
solution was l e f t standing with an excess of barium hydroxide,
the amount of barium carbonate p r e c i p i t a t e d increased slowly
with time.
In the case of methylamine, b o i l i n g with barium
hydroxide resulted i n a 100% y i e l d of barium carbonate.
These r e s u l t s suggest that the product i n question i s a carbamate
or s i m i l a r compound, which i s known to form by reaction of
carbon dioxide and amines or. .ammonia under moderately basic
condition, and which i s decomposed by a c i d or by strong base.
However, attempts to isolate, and characterize t h i s product
were unsuccessful and some question as to i t s i d e n t i t y remains.
The representation of the reaction products by equation (3-1)
i s thus subject to q u a l i f i c a t i o n , i n certain cases, although
the reactant stoichiometry appears to apply i n every case.
18
TABLE 3
RESULTS OF STOICHIOMETRY MEASUREMENTS (III)
Amine
I n i t i a l Amine PerS i l v e r Carbonate Carbonate
(Ag(l)) chlorate Amine Reduced Produced
Yield
mole/1
mole/1 mole/1 mole/1
mole/1 CO^/^Ag
Triethylamine
0.0374
0.1
0.9
0.0356
0.0179
99%
Diethylamine
0.0314
0.100
0.386
0.0281
0.0136
97%
Ammonia
0.0300
0.00
0.300
0.0172
0.0083
97%
Methylamine
0.0500
0.00
0.300
0.0455
0.0069
31%
Ethylamine
0.0400
0.00
0.300
0.0317
0.0053
33%
Ethanolamine
0.0500
0.00
0.300
0.0301
0.0086
57%
Dlethanolamine
0.0500
0.00
0.300
0.0339
0.0094
55%
Triethanolamine
0.0304
0.010
0.100
0.0243
0.0087
72%
Ethylenediamine
0.0114
0.00
0.200
o.oiii
0.001
20%
1,3-Diaminopropane
0.0400
0.00
0.300
0.0345
0.0084
49%
1
CO uptake was 0.0105 mole l "
1
1
(86%).
KINETICS AND MECHANISM
Among the s i l v e r amine complexes which were examined i n
this study, a number, including the complexes of ethylamine,
methylamine, diethylamine, ethanolamine and dlethanolamine
exhibited very similar behaviour (designated as "standard")
and w i l l be discussed f i r s t .
The triethylamine-, triethanolamine-
19
and certain diamine- complexes, exhibited some departures from
t h i s "standard" behaviour and w i l l be considered l a t e r .
ETHYLAMINE COMPLEX
This system
9
t y p i c a l of those e x h i b i t i n g "standard"
behaviour, w i l l be considered i n some d e t a i l .
The disappearance of Ag(I) at constant CO pressure obeyed
f i r s t order k i n e t i c s i n a l l experiments as shown by the t y p i c a l
f i r s t order plots of l o g [Ag(I)) v s . time i n Figure I I I . This
was v e r i f i e d by the fact that the same rate constant was obtained f o r two d i f f e r e n t i n i t i a l s i l v e r i o n concentrations
keeping the other conditions unchanged (Experiments l c and l g ) .
The rate-law obeyed during the course of each experiment
i s thus
dt
(3-2)
where (Ag(I)) i s the t o t a l concentration of a l l the Ag(I) species,
t i s time i n seconds and l o g i s common logarithm.
When a l l the other conditions were kept constant at 2 5 ° C ,
k" exhibited f i r s t order dependence on the CO p a r t i a l pressure
as shown by the p l o t of k" vs. [CO) i n Figure IV. The CO
concentration i n the reaction solution was calculated using
the s o l u b i l i t y data of S e i d e l l (10), assuming that the solution
i s saturated with CO and the s o l u b i l i t y can be approximated by
that i n pure water.
Equation (3-2) can then be rewritten, as
a second order rate-law.
20
•
0
1,000
2,000
3,000
Time,
Pig. I I I .
4,000
5,000
sec.
T y p i c a l Rate P l o t s f o r Ethylamine
See Table 4 f o r experimental
Complex
conditions.
21
i
2
6
4
(CO],
10~
4
8
mole
10
-1
Pig.IY.Dependence of Rate on Carbon Monoxide Concentration
at 25° for Ethylamine Complex;
(LH )=0.1 & (L)=0.2 mole 1
+
22
- k (Ag(X)) (CO)
(3-3)
9
where
2.303
Values of the second order rate constants, k°, measured under
various conditions are summarized in Table 4.
TABLE 4
RATES OF REACTION OF ETHYLAMINE AND RELATED AMINE COMPLEXES OF
SILVER AND OF SOME OTHER SIMILAR SYSTEMS
Initial
Amine
Ethylamine
Methylamine
CO
Pressure
Amine
Ug(I)l mm.
M
MxlO" Hg.
3
Amine
Perchlorate
M
T°C.
M
sec.
Si
o
xl0~
No.
OCt
1
z
10.0
10.0
10.0
10.0
10.0
10.0
20.0
10.0
10.0
10.0
10.0
10.0
10.0
10.0
730
730
730
730
730
730
730
562
490
395
234
742
723
712
0.200
0.200
0.200
0.200
0.300
0.100
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.0500
0.0750
0.100
0.200
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
25
25
25
25
25
25
25
25
25
25
25
20
30
35
0.310
0.206
0.152
0.077
0.153
0.156
0.149
0.156
0.154
0.152
0.143
0.103
0.223
0.337
1.55
1.55
1.52
1.54
1.53
1.56
1.49
1.56
1.54
1.52
1.43
1.03
2.23
3.37
la
lb
lc
Id
le
If
lg
lh
li
lj
lk
11
lm
In
10.0
10.0
10.0
10.0
20.0
10.0
10.0
10.0
10.0
10.0
10.0
730
730
730
730
730
562
395
234
747
742
723
0.200
0.200
0.200
0.100
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.100
0.200
0.300
0.200
0.200
0.100
0.100
0.100
0.100
0.100
0.100
25
25
25
25
25
25
25
25
15
20
30
0.283
0.152
0.094
0.152
0.152
0.283
0.287
0.286
0.110
0.174
0.424
2.83
3.04
2.88
3.04
3.04
2.83
2.87
2.86
1.10
1.74
4.24
2a
2b
2c
2d
2e
2f
2g
2h
2i
2j
2k
i
23
TABLE 4 (Continued)
Initial
Amine
CO
Pressure
Amine
(Ag(I)l mm.
MxlQ~ Hg.
M
J
Amine
Perchlorate
/M
M
T°C. sec."
k
-1
exp
see."
xl0"
No.
8.3
8.7
8.1
6.9
8.7
7.6
8.5
8.9
8.7
5.9
13.2
18.3
3a
3b
3c
3d
3e
3f
3g
3h
3i
3j
3k
31
z
Diethylamine 10.0
10.0
10.0
10.0
10.0
20.0
10.0
10.0
10.0
10.0
10.0
10.0
234
234
234
234
234
234
730
562
395
742
728
712
0.200
0.200
0.200
0.400
0.100
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.100
0.200
0.300
0.100
0.100
0.100
0.200
0.200
0.200
0.200
0.200.
0.200
25
25
25
25
25
25
25
25
25
20
30
35
0.83
0.435
0.271
0.69
0.87
0.76
0.425
0.445
0.435
0.295
0.66
0.91
Ethanolamine 10.0
10.0
10.0
10.0
10.0
736
736
736
728
718
0.200
0.200
0.100
0.200
0.200
0.0500
0.0250
0.0250
0.0500
0.0500
25
25
25
30
35
0.057
0.109
0.109
0.091
0.147
0.29
0.27
0.27
0.46
0.74
4a
4b
4c
4d
4e
Diethanolamine
736
736
736
728
718
0.200
0.200
0.100
0.200
0.200
0.0250
0.0500
0.0250
0.0500
0.0500
25
25
25
30
35
0.122
0.060
0.127
0.107
0.170
0.31
0.30
0.32
0.53
0.85
5a
5b
5c
5d
5e
10.0
10.0
10.0
10.0
10.0
The rate i s seen to be independent of the free amine concentration,
but inversely dependent on the concentration of the conjugate
acid of the amine (amine perchlorate, designated as LH ) as
+
shown i n Figure V. Thus the complete rate-law i s
-
<i(Afi(I))
dt
88 k
,
exp
(Afi(I)1 foO)
(LET]
<">
3
5
where
(3-6)
24
for Ethylamine Complete j (Lj- 0.2 H, p
- 730 mmHg
R e f e r r i n g t o (3-1) and (3-5), some r e t a r d a t i o n o f the r a t e as
the r e a c t i o n proceeds
due t o i n c r e a s e o f ( L H ) i s expected.
+
U s u a l l y t h e i n i t i a l c o n c e n t r a t i o n o f amine p e r c h l o r a t e was
made s u f f i c i e n t l y h i g h so t h a t t h i s e f f e c t was n e g l i g i b l e and
good f i r s t order p l o t s were o b t a i n e d .
at
low
However, i n some cases
( L H ) , ( A g ( I ) ) v s . t p l o t s showed the expected r e t a r d a +
t i o n i n l a t e r stages o f t h e r e a c t i o n (Expt. 1-a i n F i g u r e I I I ) .
In
these cases the i n i t i a l l i n e a r p o r t i o n o f t h e r a t e p l o t s
was employed f o r the d e t e r m i n a t i o n o f the r a t e c o n s t a n t .
OTHER "STANDARD" SYSTEMS
Among those amines which were i n v e s t i g a t e d , methylamine,
d i e t h y l a m i n e , ethanolamine
and diethanolamine
showed the same
type o f k i n e t i c s as ethylamine, i . e . , f i r s t o r d e r dependence
on Ag(I) and CO, independence o f f r e e amine c o n c e n t r a t i o n and
i n v e r s e f i r s t order dependence on ammonium i o n c o n c e n t r a t i o n .
The experimental r e s u l t s expressed i n terms o f t h e second
r a t e c o n s t a n t , k' (3-3) and k
e X
p
order
(3-5) a r e summarized i n Table 4.
Most o f t h e k i n e t i c measurements were done a t 25°C.
Typical
r a t e p l o t s f o r these systems a r e shown i n F i g u r e V I .
I n some
cases the f i r s t order r a t e p l o t s e x h i b i t e d some downward
c o n c a v i t y i n the l a t e r p a r t o f t h e r e a c t i o n .
T h i s may be
a t t r i b u t a b l e t o some zero order r e a c t i o n o f s i l v e r ( I ) w i t h
a l i t t l e i m p u r i t y i n t h e amines o r amine i t s e l f , o r i t may
be due t o some heterogeneous r e a c t i o n on s i l v e r
metal.
26
0
500
1,000
1,500
2,000
2,500
Time, sec.
Pig. 71.
Typical Rate Plots f o r Ethylamine-type Complexes
See Table 4 f o r experimental conditions
27
In those c a s e s s l o p e o f the i n i t i a l linear portions of the rate
plots were employed to determine the rate constants.
Mechanism
The inverse dependence of the rate on (LH ) may be
+
understood by taking account of the following equilibria which
prevail i n the solution.
AgL,
+
J^U
AgL + t ;
+
tAfr+KL?
UgL.2; ;
*
L +
=
H0
LH
2
+
+
jj° "''
OH";
H
(3-7)
i
= 1^
(3-8)
Referring to the stability constants of silver amine complexes
which are summarized in Appendix I, i t i s seen that the silver
ions in these solutions are present predominantly as the biscomplex, AgL , so that
+
2
(Ag(I)) &
(AgL )
(3-9)
+
2
The resulting rate expression obtained from equations (3-5, 7, 8
and 9) i s
-
4<Mm . ^
^-l K ^ - l (AgL ) (OH') (CO)
(3-10)
+
This may be identified with the following reaction mechanism
AgL
2
+ H0
L-Ag-OH + LH
2
k
0
II
> L-Ag-C-OH
L-Ag-OH + CO
0
II
L-Ag-C-OH + Ag(I)
(Rapid equilibrium)
(i)
(Rate-determining step)(ii)
>• Products
(Rapid)
(iii)
The apparent rate constant of disappearance of Ag(I), k^p*
defined by equation (3-5) and the bimolecular rate constant, k,
28
of the process ( i i ) are thus related through
= 2kK = 2 1 ^ ^ ^
(3-11)
4-
where
-
i s the association constant of AgL with OH , i.e.,
LAg
4-
- k
4- OH
K
L-Ag-OH;
fL-Ae-OH}
(AgL+} (0H~)
=
\
( 3
"
1 2 )
The factor of 2 reflects the fact that «ach rate-determining
reaction results i n the reduction of two silver ions.
Hence,
i t i s more appropriate to express the rate of the reaction i n
terms of the rate of consumption of CO, i.e., the rate of the
rate-determining
~
i L
step.
Thus,
d | = " % dt
1
d (
( I >
^ = k (L-Ag-OH) (CO)
(3-13,a)
= k^fAgL ") (OH"] (CO)
(3-13,b)
• W d ,
<- '>
4
3
13 c
Processes (3-7), (3-8) and (3-12), which are involved i n
equilibrium (i) or the process ( i ) i t s e l f i s presumably s u f f i c i ently rapid that (i) can be regarded as a pre-equilibrium.
The overall stoichiometry requires that the reaction intermediate
containing a CO molecule, and one silver ion, L-Ag-COOH, reacts
with another silver (I) species (the identity of which i s to be
discussed later).
However, this step ( i i i ) appears to be fast,
compared with ( i i ) , so that the kinetics are f i r s t order in Ag(I).
0
This intermediate complex, L-Ag-C-OH, i s analogous to the one
29
which was previously proposed by Harkness and Halpern (3) as an
intermediate complex in the reaction of Hg , i.e., -Hg-C^-OH.
2+
Support for the structure of the latter was provided by the
observation by Halpern and Kettle (6) that when methanolic
solution of mercuric acetate takes up CO under similar conditions,
a stable methylformate derivative, AcO-Hg-H-OCH^, analogous to
the proposed complex was formed, isolated and spectroscopically
identified.
They also reported that attempts to prepare analogous
CO adduct of silver acetate were unsuccessful, but this can presumably be attributed to the poor solubility of silver acetate
in
methanol and instability of the CO adduct toward decomposi-
tion into metallic silver.
In terms of this mechanism i t might be expected that the
rate constants k (and also kK ) should be relatively insensitive
n
to the nature of L and thus that the large dependence of k^^ on
the nature of the amine should reflect largely the variation of
and K^.
This i s shown to be the case in Table 5 where i t
is seen that notwithstanding a 30-fold variation in k^p
-3
-2
-1
(which ranges from 2.8x10
to 8.6x10
sec. , for the five
amines under consideration) the value of kK^ i s substantially
constant, (1x10
5
mole
-2
2
-1
1. sec. ) for a l l the systems.
30
TABLE 5
SUMMARY OF KINETIC AND RELATED THERMODYNAMIC DATA
FOR "STANDARD" SYSTEM AT 25°C.
2)
1)
lc _____
exp
Amine
C H NH
xlO
mole
l.-l
%
h
xlO
4
4
sec."*"
3)
kK
mole.
l."
1
xlO"_ 9
mole 1?
sec."
5
xlO*
mole
I"?
1.55xl0"
2
1.2
6.5
2.85xl0"
2
2.9
5.2
15
0.9
(C H ) NH
8.6 xlO"
2
5.0
9.1
46
1.0
HOC H NH
2.8 xlO"
3
2.8
0.55
1.5
1.0
(HOC H ) NH
3.1 xlO"
3
0.10
1.6
1.0
2
5
2
CH NH
3
2
2
5
2
2
4
3
2
4
2
16
7.6
1.0
1) and 2) Refer to Appendix I.
3) Calculated by use of equation (3-11).
The temperature dependence of. the rate constants was
determined for a l l these systems over the temperature range
15 to 35°C.
In a l l cases good linear Arrhenius plots were
obtained, which are given in Figure VII.
parameters are summarized i n Table 6.
The activation
31
3.2
3.4
3.3
T" ,
1
Fig. VII.
10~
3
3.5
dee'
1
Arrhenius Plots f o r Ethylamine-type Complexes
32
TABLE 6
APPARENT ENTHALPY AND ENTROPY OF ACTIVATION
FOR "STANDARD" SYSTEMS
*
1)
3
AS
A H exp
exp
Kcal. mole **e.u.mole-1
exp
Amine
C2H-NH2
7.8 xlO"
3
14.3
-20.0
CH_NH
1.43xl0"
2
15.5
-14.9
(C H-) NH
4.3 xlO"
2
14.1
-17.5
HOC H NH
1.4 xlO'
3
17.4
-13.1
(HOC H ) NH
1.6 xlO"
3
18.3
- 9.9
2
2
2
2
4
2
1)'
2
4
2
%k
,> i s used for the calculation of the activa- exp = kK hKjd^KT
u
tion parameters.
The present result shows that the only silver (I) species
that i s active toward CO under the conditions investigated
for these five amines i s the hydrolyzed mono-amine complex,
L-Ag-OH while other silver (I) species involving the bis-complex,
AgL2 , and free silver ion or aquo complex, Ag , make a negligible
+
+
contribution.
Although metallic silver was precipitated during the
course of reaction, i t s heterogeneous catalytic effect, at
least during the early stages, was small.
33
TRIETHYLAMINE COMPLEX
The triethylamine complex exhibited somewhat different
kinetic behaviour, from the "standard" systems described above,
notably in that the rate of reaction was no longer independent
of the free amine concentration but exhibited an inverse dependence on the latter.
The rate was inversely proportional to
ammonium concentration but the effect of silver concentration
was somewhat more complicated.
The experimental results are
summarized in Table 7 and typical rate plots are shown i n
Figure VIII.
This complex exhibited the fastest overall rate of a l l
the amine complexes investigated in this study, so that a
32.3% CO - 67.7% N mixture was used in most of the experiments
2
to obtain a reaction rate convenient for measurement. A f a i r l y
high concentration of triethylamine (up to 0.8 mole'l ^) was
employed to prevent the hydrolysis of silver and precipitation
of silver oxide because the triethylamine-silver complex i s
much less stable than the complexes of primary and secondary
amines, while the basicity of the amine i s almost the same.
Because of the high vapor pressures of the resulting solutions
Lattey's (8) data on the vapor pressure of aqueous triethylamine
solution were used to calculate the partial pressure of CO.
34
T A B L E
R A T E
O F
R E A C T I O N
O F
T R I E T H Y L A M I N E
1)
Amine Vapor
Initial
Amine
Per- Pres(L)
chlorate sure
(Ag(I)h
mole-l" - mole-l" mmHg
mole°l
1
O o O l O
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.005
0.005
0.005
1)
0.809
0.802
0.794
0.784
0.205
0.250
0.263
0.309
0.342
0.412
0.508
0.619
0.080
0.180
0.180
1
0.600
0.480
0.400
0.300
0.600
0.600
0.600
0.600
0.600
0.600
0.600
0.600
0.300
0.300
0.200
7
79
79
78
78
47
51
53
57
60
65
70
74
34
45
45
C O M P L E X
A T
2 5 ° C .
C O
PresK
sure mole'l
mmHg sec"*220
220
220
220
230
229
228
227
226
224
223
222
234
231
231
1.02
1.31
1.71
2.27
2.32
2.16
2.12
1.91
1.74
1.62
1.45
1.20
5.1
3.4
4.9
,
k
exp
sec"*-
Expt.
No.
0.61
0.63
0.68
0.68
1.39
1.30
1.27
1.15
1.04
0.97
0.87
0.72
1.53
1.02
0.98
6a
6b
6c
6d
6e
6f
6g
6h
6i
6j
6k
61
6m
6n
6o
o v n
Reference (8).
Although triethylamine was very carefully purified as
described i n the experimental section, there was an appreciable
amount of i n i t i a l reduction of silver (up to 25% of total
silver (I) concentration) which may be attributed to some
impurity
i n the amine, and which introduced some complications,
* The reaction mixture was usually stable toward the i n i t i a l
reduction under nitrogen flow especially at high ammonium concentration. This i n i t i a l reduction occurred usually after C O
gas was introduced into the reaction vessel, as can be, seen i n
the rate plots (Figure V I I I ) .
The amount of the i n i t i a l reduction of silver was directly dependent: on the total amine concentration (free amine and ammonium) i n the reaction solution.
However, i t seemed to have no significant effect on the rate of
the subsequent reduction by C O .
35
36
particularly in the determination of the effect of i n i t i a l silver
concentration on the reaction rate, since at low i n i t i a l (Ag(I))
a large portion of Ag(I) was lost by the. i n i t i a l reduction
while high (Ag(I)J could not be employed because of the instab i l i t y of the complex toward hydrolysis of silver ion and precipitation of silver oxide.
Two alternative interpretations of the inverse dependence
of the rate on the free amine concentration were considered.
The f i r s t of these involved the possibility that not only the
mono-complexed silver ion but also uncomplexed ion (or aquo ion)
contributes to the overall reaction to an appreciable extent.
This failed to yield a rate-law which f i t t e d the observed
dependence of the rate on the silver (I) and amine concentrations
in detail, and, furthermore, required that the reactivity of
free silver ion, which i s present i n very low concentration
compared to the mono-complex, be much higher (some 300-fold)
than that of the latter.
This could not be readily reconciled
with the observed insensitivity of the rate of various other
mono-complexes to the nature of the amines, and with the measurements of Peters and McAndrew (5) on the reduction of uncomplexed
Ag
+
ions by CO i n perchlorate media.
Consequently this inter-
pretation was rejected.
A more satisfactory account of the. kinetic behaviour of
this system was obtained in terms of the following mechanism i n
which the competitive reaction of the intermediate, L-Ag-COOH,
37
are considered.
k
L-Ag-OH + CO
x
v
k
+
k
L
- L-Ag-COOH
(ii')
-l
2
L=Ag-COOH + LA.g
> Products
Here, processes k ^ and
(iv)
are competing for the intermediate
complex, L-Ag-COOH, the back reaction of the f i r s t step ( i i ) no
longer being negligible.
Assuming steady state approximation
for L=Ag-COOH, the kinetics are found to be
_ d(CO)
, dCAn(I)) _ .
tAg(I)] (CO)
dt
*
dt
exp
(LH+)
m
%
.
k
k lW
k
r
i
l h b d
K
K
,
Vh
k
k
exp
"
Thus, k
K
K
k.j+kj^
(LH*")
K
L
(A (I)j (CO)
VVSi,
(LH+)
kK
K
2
k~ -1CL)+k.I.Kd^ (Ag(I)j
•L
1
t
.
l
t
i s no longer a constant.
ktp" " 2 k
l
V
k
2
K
>
d
(Ag(I))
CAg(D)/(L)
1
2
fi
^iVW^** "
2VAg(I))/(L)
L W y
*
3 14
>
,~ , .
^- >
(Ag(I))
3
(3
15
" >
16
Taking reciprocals
fr^CL^K
(Ag(I)))
(3-17)
l
From the equation (3-25), we expect a linear relation between
1/k^p
and ( L ) at a fixed total silver ion concentration,
(Ag(I)) . The plots of l/k „ vs. ( L ) i n Figure IX based on
ov
the experimental results in Table.7 are in complete accord with
this.
From equation (3-17) the intercept and the slope of the
linear plots are given by
3 8
0.0
0
0.2
0.4
0.6
(L),
Pig. IX.
mole l "
0.8
1
Dependence of Rate on Free Amine Concentration
at 25° f o r Triethylamine Complex
39
Intercept - fc g v
1 n D d.
2
A „
Slope =
(3-18)
K
Hr „ r L / 0
1
(3-19)
T
k
2
K
d
[ A 8 ( I ) )
1
Thus, the intercept should be independent of (Ag(I)) and the
slope should be inversely proportional to (Ag(I))
s
i n accord
with Figure IX, which also shows k. _ to be independent of (LH }.
+
From the intercept and slopes of the plots i n Figure IX, and
using equation (3-18) and (3-19), the following values were
obtained for the rate constants.
k - ^ - 2.5xl0 mole" «1 .sec."
5
k
2^ -l
k
=
3
'
4 x l
°
3
2
2
(3-20)
1
mole" '!
1
(3-21)
The value of kjK^ i n this case i s about 2.5 times as large as
that found for ethylamine and related complexes.
It i s obvious that when the second term i n the denominator
of the equation (3-16) i s much larger than the f i r s t term,
(i.e., k^K^ (Ag(I))^ k_^ L ), the overall kinetics approach
those for the ethylamine complex.
For some reason, i n the case
of triethylamine, these two terms are comparable i n their magnitudes.
Since
for triethylamine i s actually larger (70 times)
than for ethylamine, the observed kinetics must reflect either
an abnormally large value of k_^ or an abnormally small value
of k£.
40
TRIETHANOLAMXNE COMPLEX
The only other tertiary amine that was investigated i n this
study was triethanolamine.
The silver complex of this amine
exhibited almost the same kinetics as ethylamine, the rate i n
this case being almost independent of the free amine concentration;, with only a slight dependence i n the opposite direction to
that for triethylamine (i.e., the rate increasing with amine
concentration) and inversely proportional to the ammonium concentration.
The experimental results are summarized i n Table 8
and typical rate plots are given i n Figure X.
The dissociation
constant of the triethanolamine complexes i s so large (K
that It i s no longer valid to approximate [AgL
2
d
= 0.046)
) by (Ag(I)) .
This would give rise to a dependence of the rate on the free
amine concentration even i f the reaction followed the same
kinetics as for the ethylamine complex.
The observed dependence
of the rate on the amine concentration i s i n the direction
expected from this (i.e., the rate increases with the amine
concentration), but i s much smaller than predicted.
This suggests
that there i s also superimposed upon this an inverse dependence
of the rate on the free amine concentration, similar to that
found for triethylamine complex.
This i s not unexpected i n
* Actually, using the value for K^., given above, the calculation shows that more than 30% of the silver i s present i n the
form of AgL at (L) = 0.1 mole«l ^
+
m
0
41
TABLE 8
RATE OF REACTION OF TRIETHANOLAMINE COMPLEX
Initial
(Ag(I)1
mole*I"
*
1
Amine
CO
k«
Amine
PerPresL
chlorate sure
mole°*l»l
mo 1 e • 1" ^mo 1 e • 1 °°n1
nnHg T°C. sec'l.
k
exp
sec'l
Expt.
No.
0.010
0.100
0.0500
705
40
0.23
1.2x10'-2
7a
0.010
0.100
0.100
705
40
0.13
1.3x10'-2
7b
0.007
0.100
0.200
705
40
0.069
1.4x10'-2
7c
0.07
0.100
0.100
705
40
0.13
1.3x10'-2
7d
0.007
0.200
0.100
705
40
0.13
1.3x10"-2
7e
0.007
0.040* 0.0500
730
25
0.049
2.4x10"-3
7f
0.007
0.100* 0.0500
730
25
0.060
3.1x10'-3
7g
0.007
0.190* 0.0500
730
25
0.062
3.1x10'-3
7h
0.014
0.040* 0.0500
730
25
0.051
2.6x10"-3
7i
0.007
0.040*
0.0250
730
25
0.091
2.3x10'-3
7j
0.007
0.100
0.0500
728
30
0.106
5.3x10'-3
7k
0.007
0.100
0.0500
718
35
0.17
8.5x10'-3
71
Free amine concentrations corrected for the dissociation of
silver complex.
view of the similar stability constants of the two complexes.
The dissociation constants of both complexes are much larger
than that of the ethylamine complex and i n both cases the
order of the f i r s t and second dissociation constants ( K ^ ^
is the reverse of that for ethylamine and other primary and
)
42
-1.8
Log(Ag(l)]
0
1,000
2,000
3,000
Time,
Fig. X.
- i —
4,000
5,000
sec.
Typical Bate Plots f o r Triethanolamine Complex
See Table 8 f o r experimental conditions.
43
secondary amines (Appendix I ) .
Applying the same mechanism to
triethanolamine, as previously to triethylamine, and assuming
(Ag(I) ) = U g L ) + [AgL )
+
(3-22)
+
2
(3-23)
K
+ CD
d
d
l
and neglecting other silver species, then the rate law i s
expected to be of the form
,
*
^
dlCOJ _
dt
d(An(I>]
.
dt
.
(Afi(I)_j (CO)
* ""exp
ex
fur)
^
k
P
k (AgL )
+
2
k
( A g L ) (OH") (CO) ^J^^
(3-24)
+
lKh
2
fD
k ^ ( K +(Lj)^fk K (Ag(I)}(K +(L))
1
1
1
(3-25)
s
d
2
d
d
It was not possible to make sufficiently accurate kinetic
measurements to test this equation i n detail and determine a l l
the rate constants involved.
However, an attempt was made
to estimate the rate constant k^ from the experimental rate
constants.
It i s seen from equation (3-25) that when ( D
i s very small,
the f i r s t term in the denominator becomes negligible compared
to the second term and the kinetics approach the following form.
d[CO
dt
=
k
l h b d
K
K
K
fAe
1
CLH^ ) °'
f
C
1
K, +(L)
f
(3-26)
44
Among the experimental data Expt. 7-f i s the one done at lowest
( L ) (=0.040 mole*!"*-).
For the experimental condition of this
experiment the ratio of the two terms in the denominator of
(3-25) was calculated to be
-l
k
k
2
( d j + (IQ)
K^(Ag<I))(K +(L])
K
2
d
using the value of
" °'
08
given in Appendix I (K^
1
( 3
"
2 7 )
= 0.046) and
1
assuming the same value of k.^/k^ as for triethylamine given
by (3-21).
This ratio seems to be sufficiently small to approxi-
mate the kinetics at these experimental conditions of Expt. 7-f
by (3-26).
(Simple comparisons of the numerical values of
( L ) / ( K , + ( L ) ) with these experimental results in Table 8 show
that below ( L ) = 0.1, the results agree f a i r l y well with the
kinetics given by (3-26) but deviate from i t very rapidly as
( L ) increases. This i s expected since the f i r s t term i n the
denominator of (3-25), which was neglected i n (3-26), i s second
order i n (L).) From (3-26) k ^
k-K. =
kexp
i s given by (3-28).
d
l
(3-28)
Using the data of the Expt. 7-f, k^K^ was estimated to be
1x10
5
=2
mole *" 1
2
sec
-1
. Although this i s a very rough estimate,
i t agrees with that of ethylamine (kK^ = 1.0x10"*), at least in
order of magnitude.
A further expected feature of the kinetics in the observed
45
curvature of Arrhenius plot (Figure XI) of the apparent rate constant, k
j at ( L j = 0.100 mole-l" . As can be seen from the
* exp'
1
several experiments at 40°C. (Table 8) the kinetics at this
temperature were closer to the "standard" ones than at 25°C.
However, at (L) - 0.1 and at 25°C. the kinetics appear to be
close to those for ethylamine and the two other ethanolamine
complexes and the apparent energy of activation under this condition and at this temperature i s also very close to those for two
*
-1
other ethanolamines ( AE
~
18 Kcal. mole ).
This mechanism, comprising a sequence of ( i i ) and (iv),
1
appears to give a satisfactory account of the kinetic behaviour
of the triethylamine and triethanolamine complexes, although only
a semi-quantitative discussion was possible for the latter.
As
mentioned previously, the necessary condition for this mechanism
to hold i s that either k_^ i s abnormally large or k
small compared to those for ethylamine.
2
abnormally
This behaviour may be
related to the fact that the dissociation constants (especially
K^) of the silver complex of tertiary amines are abnormally
large (K^ i s 70 times for triethylamine and 400 times for
triethanolamine complex as large as that of ethylamine; Appendix
I ) , that i s , their bis-complexes seem to be abnormally unstable
toward the loss of the second coordinated group.
A similar
instability of the intermediate complex L-Ag-COOH toward decomposition could account for an abnormally large value of k_^.
These observations provide some information about the
46
-1.0
Ethylene-,! amine
ChH ] - 0.100 mole 1"
+
(L) - 0.200 mole 1*
-1.5
Log k
exp
-2.0
Trie thanoland ne
(LH J - 0.050
+
(L ) - 0.100 mole 1
-2.5
3.2
3.1
3.3
T" ,
1
Pig. XI.
10"
3.4
3
degT
1
Arrheniue Plots f o r Triethanolamine
and Ethylenediamine Complexes
3.5
47
second step of the reaction, beyond that which could be deduced
from the results of the ethylamine-type systems„
In particular,
they indicate that the second silver species which reacts with
the intermediate complex i s also a mono-complexed species.
This point i s to be discussed again later.
AMMONIA COMPLEX
The ammonia complex exhibited the slowest overall reaction
rate of a l l the amine complexes investigated in this study.
Most of the rate measurements for this system were made in an
autoclave at 30°C. and at high CO pressure (up to 20 atm.).
The experimental results are summarized in Table 9 and some
typical rate plots are given in Figure XII.
The dependence of
the rate on the ammonium ion concentration was complex. The
log k' vs. log (LH ) plot i n Figure XIII shows that the rate
+
is inversely proportional to (LH ) at low ammonium ion concentrations but that at higher ammonium ion concentrations the inverse
order increases to two or higher.
This implies that the reaction
mechanism may be different i n the two regions or the rate i s
controlled by different steps.
The effects of silver ion concentration, CO pressure and
free ammonia concentration were studied at both low (0.02
mole'l"''") and high (0.1 mole»l~*) NH^ concentration and these
+
results also are summarized in Table 9. In both NH^
region,
the reaction rate was f i r s t order i n CO and only slightly
48
TABLE 9
RATE OF REACTION OF AMMONIA COMPLEX
Initial
CO
Ammonia Ammonium
k*
(Ag(I);) Pressure CU
(LR+)
mole ^»l
mole°l
atm. mole-1" mole•1~1 T°C. sec'^=
exp
sec 1
Expt,
No.
Effect of Ammonium
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
9.2
9.2
9.2
9.2
9.2
9.2
9.2
9.2
9.2
9.2
9.2
0.180
0.180
0.180
0.180
0.180
0.180
0.180
0.180
0.180
0.180
0.180
0.0200
0.0236
0.0300
0.0500
0.0700
0.0850
0.100
0.125
0.150
0.175
0.200
30
30
30
30
30
30
30
30
30
30
30
4.6x10 -2
3.3x10-2
2.8x10-2
1.65x10-2
7.6x10-3
8.3x10 -3
5.5x10 -3
5.0x10 -3
3.1x10 -3
1.8x10-3
1.2x10-3
9.2x10-4
7.4x10 »4
8.4x10 -4
8.3x10 -4
5.3x10 -4
7.0x10 -4
5.5x10 -4
6.3x10-4
4.7x10 -4
3.1x10 -4
2.4x10-4
8a
8b
8c
8d
8e
8f
8g
8h
81
8j
8k
Effect of CO Pressure
0.010
0.010
0.010
0.010
4.1
9.2
14.1
19.2
0.180
0.180
0.180
0.180
0.020
0.020
0.020
0.020
30
30
30
30
5.3x10-2 10.5x10-4
4.6x10 -2 9.2x10-4
4.9x10 -2 9.8x10-4
5.0x10 -2 10.0x10 -4
81
8a
8n
8o
0.010
0.010
0.010
0.010
4.1
9.2
14.1
19.2
0.180
0.180
0.180
0.180
0.100
0.100
30
30
30
30
5.7x10-3
5.5x10-3
5.7x10-3
7.2x10-3
5.7x10-4
5.5x10 -4
5.7x10 -4
7.2x10-4
81'
8g
8n'
8o»
o.ido
0.100
Effect of I n i t i a l Silver Ion
9.2
9.2
9.2
0.180
0.180
0.180
0.020
0.020
0.020
30
30
30
4.2x10'-2
4.6x10'-2
4.8x10'-2
8.5x10'-4
9.2x10"-4
9.6x10'-4
8p
8a
8q
0.0048 9.2
0.0099 9.2
0.0196 9.2
0.180
0.180
0.180
0.100
0.100
0.100
30
30
30
2.6x10 -3
5.5x10'-3
9.8x10'-3
2.6x10"-4
5.5x10"-4
9.8x10"-4
8p'
8g
8q»
0.0047
0.0047
0.0195
49
TABLE 9 (Continued)
Initial
CO
Ammonia Ammonium
(Ag(I)} Pressure (L j
(LH+)
mole•I" atm. mole-l mole !" T°C.
1
1
4
1
k°
mole" -!
sec"l
k
sec"
1
e
x
Expt
No.
p
1
Effect of Free Ammonia
0.010
0.010
0.010
9.2
9.2
9.2
0.080
0.180
0.280
0.020
0.020
0.020
30
30
30
3.9x10-2
4.6x10-2
3.1x10-2
7.8x10-4
9.2x10>4
6.2x10-4
8r
8a
8s
0.010
0.010
0.010
9.2
9.2
9.2
0.080
0.180
0.280
0.030
0.030
0.030
30
30
30
2.7x10-2
2.8x10-2
2.2x10-2
8.2x10=4
8.4x10-4
6.6x10-4
8r"
8c
8s"
0.010
0.010
0.010
9.2
9.2
9.2
0.080
0.180
0.280
0.100
0.100
0.100
30
30
30
7.6xl05.5x10-3
5.1x10*3
7.6x10-4
5.5x10"
5.1x10-4
8r
8g
8s»
0.050
0.050
0.050
0.050
0.050
25
30
35
40
50
1.06x10-2
1.4x10-2
2.0xl0~
3.1x10*2
6.6x10-2
5.3xl0"
7.2x10-4
10.2x10-3
1.5x10-3
3.3xl0"
8t
8d
8u
8v
8w
3
4
e
Effect of Temperature
0.010
0.010
0.010
0.010
0.010
9.2
9.2
9.2
9.2
9.2
0.180
0.180
0.180
0.180
0.180
dependent on the free NH^ concentration.
2
4
3
No trend, however, was
discernible i n the latter dependence. The dependence on
(Ag(I)j was f i r s t order at low (NH^ ) and second order at high
+
(NH ) (i.e., k' was proportional to (Ag(I)).)
+
4
This suggests that
at high (NH^ ) the rate i s controlled i n part by the second step
in which the second silver (I) species takes part.
The high inverse
order dependence on [ NH^ J i n this region further suggests that
the second silver species also is hydrolyzed.
50
Time,
Pig. XII.
sec.
Topical Rate P l o t s f o r Ammonia Complex
See Table 9 f o r experimental c o n d i t i o n s .
51
-1.0
1.5
-l.o
-0.5
Log [LH ]
+
Pig. XIII.
Dependence of Rate on Ammonium Ion Concentration
at
30° f o r Ammonia Complex; [ L j = 0.200 mole l " "
1
52
The following mechanism i s consistent with these observations t
l
k
.
L-Ag-OH + CO ^
L-Ag-COOH
(ii')
-1
k
2
L-Ag-COOH + L-Ag-OH
> 2Ag + C0 + 2L + H_0
2
(v)
where two processes with rate constants k_^ and k2 are competing
for intermediate complex, L-Ag-COOH. The kinetics, assuming
steady state concentration of L-Ag-COOH, are thus of the following type.
" " ^ I P
- "^T
1
k
"
1
CAg(X)) (CO)/(LH )
+
k k K^Kg4 (Ag(I))2(C0)
=
1
2
1
" (LHT)^k^^K^K^(Ag(I))/tLH*")}
(3-29)
At low (NH^ where k_ « k ^ I ^ K ^ (Ag(I))/[LH+), the kinetics
x
approach those for ethylamine while at high (NH^ ) where
4
k_^»
k2Kj K K ^(Ag(I))/(LH
i
b
], the reaction i s second order on
d
(Ag(I)) and inverse second order on (LH J, i . e . ,
+
L o w
( L H
+)
d
8
. , +,
High (LH )
TW
s
M
m
V
d(00)
=
-
d
t
h
V
JAig^fii
^
( 3
2 2 2 (A (I)) fCOj
K ^ K ^
(LH+J-
. )
3 0
2
fi
k
-
1
(
3 _ 3 1
>
From the experiments at low (LH ) (0.020 mole l " )
+
k
exp -
Using the values for
2 k
lW
1
=
d l
9 X 1 0
'
4
SeC
"
X
and K ^ at 30°C. in Appendix I,
d
( 3 _ 3 2 )
53
(3-33)
On the other hand at high (LH*")
k
k
exp - ^1
e X
p corresponds tos
K ^ K ^ JA^IJ
The experimental rate constant
( 3
„
3 4 )
at (LH ; = 0.100 mole"1
(Expt s. 8p°, 8g, 8q') exhibits a good f i r s t order dependence on
9
(Ag(I)) and gives a constant value for k
/(Ag(I)) (5.4, 5.6
-9
-1
-1
and 5.0x10 mole •l-sec , respectively), i.e.,
A
x
at (LH ) = 0.100 mole'l"
+
2 2 2
2k, k h bKdT
k
K
( A i ^ r
Substituting
values for
•
"WT
•
5
X
1
0
2
,
9
- l - -
- ! - ^ "
1
0-35)
1
(LH J = 0.100 mole l * * and using (3-32) and the
+
1
and
again we get
2 h
k -1i
k
All
_
K
9
1
K
2xl0 mole" .!
9
these results are for 30°C.
2
(3-36)
2
This temperature was chosen
for the most of kinetic studies instead of 25°C. because of
d i f f i c u l t y in controlling the autoclave temperature at 25°C.
To compare the rate constant with those of other systems, i t i s
necessary to correct i t to 25°C. Arrhenius plot at (LH ) •
+
0.050 mole 1 * given in Figure XIV also reflects the complex
nature of the kinetics of this system.
It was d i f f i c u l t to
measure the temperature effect at lower (LH*) than this because
54
55
of the fast rates and poor rate plots. The Arrhenius plot i n
Figure XIV i s linear at high temperature but i s concave upwards
at low temperatures.
This indicates that the linear portion
represents the temperature dependence of the apparent rate
constant of the form given by (3-34) and not (3-32).
The appa-
rent energy of activation calculated for linear portion of the
Arrhenius plot i s 13.7 Kcal. mole \
which i s very close to the
values for k
(^kH^K^R^ ) for other systems (14-18 Kcal.
mole~1 , Table 6 ) . It thus ^ seems reasonable to calculate the
rate constant given by (3-32) for 25°C. using a temperature co*
»i
efficient corresponding to AH
«= 14-18 Kcal. mole . . Because
the temperature interval involved i s small (5°C), the error
involved in the extrapolation can hardly be very large.
Thus
s
from (3-32)
exp
- 210^1^^ - 6xl0~
Again, using the values for K
Vh
"- «* °
1
x l
3
sec"
1
(at 25°C.)
(3-37)
and K, for 25 C. in Appendix I.
b
°i
5
mole' »1 -sec"
2
2
1
(3-38)
This i s very close to the values previously found for ethylamine
and related systems (1.0x10^).
At (NH^ + ) exceeding 0.1 mole'l -1 , the dependence of the
rate on (NH^ ) seems to exceed inverse second order (Figure XIII).
+
However, in this region the actual rate of the reaction was
extremely slow so that the measurements are considered unreliable.
The mechanism involved here i s substantially the same as that
derived previously for the two tertiary amines«, triethylamine and
triethanolamine apart from a difference in the nature of the
9
Ag(I) species participating in the second step of the reaction
(L-Ag-OH and LAg , respectively).
difference i s available.
No explanation for this
Recently, Peters and McAndrew (5)
reported the kinetics of the reduction of silver in perchlorate
media at 70°C. to be
(3-39)
which was interpreted by the following mechanism
Ag + CO + H 0
+
2
AgCOOH + Ag
+
-=_±
AgCOOH + H
> 2Ag + C0 + H
o
(Rapid equilibrium)
+
+
(Rate determining)
In this case the second silver (I) species reacting with the
intermediate complex i s an unhydrolyzed ion analogous to LAg
in the tertiary amine cases.
The slowness of the second step in the case of the ammonia
complex which leads to the departure from the simple ethylaminetype kinetics may be due to the much lower basicity of ammonia
(i.e., to a smaller L-Ag-OH concentration).
However, the
failure of diethanolamine, which i s even less basic, to exhibit
the same type of kinetics as ammonia, throws some doubts on this.
DIAMINE COMPLEXES
In addition to the above monoamines, three primary diamines,
ethylenediamine, 1,3-diaminopropane and 1,4-diaminobutane were
examined. The experimental results are summarized in Table 10
57
TABLE 10
RATE OF REACTION OF DIAMINE COMPLEXES
Initial
CO
Amine
(Ag(I)) Pres- Amine
Persure (L) chlorate
mole" .! exp
mole.I" mmHg moleX mole l- T°C. sec™ sec".
8
k
k
1
Amine
1
1
o
1
1
1
Ethylenediamine
0.013
0.013
0.013
0.012
0.020
0.020
0.020
0.010
0.010
0.010
0.010
0.010
0.010
0.010
0.010
705
705
705
705
705
508
239
741
742
730
728
712
699
683
664
0.300
0.300
0.100
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.200
0.100
0.300
0.100
0.200
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
0.100
40
40
40
40
40
40
40
15
20
25
30
35
40
45
50
1,3-Diaminopropane
0.010
0.010
0.010
0.005
0.010
0.010
0.010
0.010
0.010
736
736
736
736
736
736
736
736
736
0.183
0.183
0.183
0.192
0.0356
0.0643
0.0840
0.183
0.282
0.100
0.050
0.025
0.100
0.025
0.025
0.025
0.025
0.025
25 7.1xl0- 7.1x!0" 10a
25 1.6X10" 8.2x10-3 10b
25 2.8x!0 7.1xl0" 10c
25 7.1x10-2 7.1x10-3 10d
25 1.7x10-1 4.2x10-3 I0e
25 2.2X10" 5.5x10-3 I0f
25 2.5X10" 6.1x10-3 log
25 2.8x10" 7.1x10-3 lOh
25 3.0x10-1 7.6x10-3 10i
0.010
0.010
0.010
0.010
0.010
0.010
0.010
736
736
736
736
736
736
736
0.190
0.190
0.190
0.090
0.050
0.010
0.005
0.0177
0.0250
0.0354
0.0250
0.0250
0.0250
0.0250
25
25
25
25
25
1 4= Diaminebutane
s
0.34
0.112
0.34
0.17
0.34
0.34
0.34
0.132
0.193
0.216
0.220
0.228
0.354
0.480
0.668
0.034
0.034
0.034
0.034
0.034
0.034
0.034
0.0132
0.0193
0.0216
0.0220
0.0228
0.0354
0.0480
0.0668
No.
9a
9b
9c
9d
9e
9f
9g
9h
9i
9j
9k
91
9m
9n
9o
3
2
1
3
<=1
1
1
1
7.6xl0- 11a
8.1x10-3 l i b
8.3x10-3 H e
8.3x10-3 H d
8.4x10-3 l i e
25 2.6X10" 6.4x10-3 l l f
25 2.4X10" 6.0xl0°3 l i
4.3x10-1
3.2x10-1
2.3x10-1
3.3x10-1
3.4x10-1
and some typical rate plots are given i n Figure XV.
3
1
1
g
The ethylene-
diamine complex was f i r s t investigated at 40°C. and found to
58
0
1,000
2,000
Time,
Pig. XV.
3,000
sec.
Typical Rate Plots f o r Diamine Coplexes
••• See Table 10 for experimental conditions.
59
exhibit kinetics similar to those of the ethylamine complex as
can be seen from Table 10.
However, the study of the temperature
dependence of the rate revealed more complicated behaviour at
lower temperature and the results failed to yield a linear
Arrhenius plot (Figure X V I ) .
Attempts to elucidate the kinetics
at 25°C. and obtain directly the rate constant at this temperature for comparison with the other amine complexes were unsuccessful.
Thus the rate of this system at this
temperature
exhibited a very complicated dependence on amine, amine perchlorate and silver ion concentration which was not altogether
reproducible.
At higher temperature, however, (above 35°C.)
the system appeared to be well behaved.
Extrapolation of the linear portion of the Arrhenius plot
in this region yielded a value of 1.07x10 * sec
at 25°C.
Using the values for
and
for
of ethylenediamine
in Appendix I , we get
kK. = exp
k
2 K
b
= 0.30xl0 mole' .l .sec"
5
2
2
1
(3-40)
K
d ] L
This i s only about one-third of the value for ethylamine
and related amine complexes.
The silver complex of 1,3-diaminopropane has an abnormally
large f i r s t stability constant (KQ-J_
88
8.9x10"*) as can be seen
in Appendix I , and there has not been reported any data on i t s
second stability constant.
This i s due to stabilization of
61
the mono-complex by chelation, with the r e s u l t that there i s
l i t t l e tendency to add a second amine molecule to form the b i s complex.
This great s t a b i l i t y of the monochelate complex i s
r e f l e c t e d i n the extremely large f i r s t s t a b i l i t y constants the
second s t a b i l i t y constant,
so that the approximation
i n v a l i d i n t h i s case.
9
presumably being very s m a l l
[Ag(I)) = CAgL^J i s obviously
Therefore, equation
(3-41)
(as i n the
case of triethanolaminej, which also has a small K^£»
K ^)
d
s
must be used for the concentration of
AgL
(AgL ) + ( A g L ) = (Ag(I))
+
+
2
+
.
i . e . , large
Assuming
(3-22)
[AgL^)
(AgU*" )(L)
=
K
12
=
K
where K i s the second s t a b i l i t y constant of the s i l v e r complex.
Then i f ethylamine-type behaviour applies also i n t h i s case,
the o v e r a l l k i n e t i c s w i l l be as follows:
COl _
dt
=
k
*
JjA^jIlL
dt
"Vb
_ fgxp_ C A R ( I ) ) (COj
2
(1+KCL))(LH+)
TLH+I
( 3
4 1 )
Then
k
exp
85 2 k K
(3-42)
h b 1+K(L)
K
Taking the r e c i p r o c a l of
k^^
>4r = 2 E ^ ( l T 7 « )
+
62
The experimental free amine concentration was f i r s t calculated
assuming
and on this basis the experimental
results for different (L Jwere plotted as 1/k
vs. 1/(L).
The resulting plot, which was not quite linear but slightly
concave upwards was used to obtain a rough estimate of K, which
9
in turn was used to improve the free amine concentration. This
procedure was repeated to self-consistency and yielded a good
linear plot of l / k p vs. l/(L) shown in Figure XVI.
eX
stants derived from this, using the value of
kK^ = 5.1xl0 mole" -l osec"
2
2
2
The con-
in Appendix I are
(3-44)
1
K - 26 mole" .!
(3-45)
1
This value of kK^ i s only 1/200 of the "normal" value for ethylamine, etc. (kK^ = 1.0x10^).
The much lower reactivity of this complex presumably reflects
blocking of the reaction site by chelation, i.e.
H N —Ag — NH
+
0
HoC -
,
H
2
CHn
2
The rate constant of ethylenediamine i s also smaller than
that of ethylamine, etc., but the reduction factor i s only 1/3
in this case.
Presumably this also i s attributable to chelation
but in this case the chelation tendency i s much smaller than
4.
1,3-diaminopropane, because of the preference of Ag
for linear
coordination which, for steric reason, i s more readily realized
with 1,3-diaminopropane than with ethylenediamine.
This i s
63
reflected also in the corresponding stability constants of the
complexes.
Thus the f i r s t stability constant of ethylenediamine
silver complex, while larger than that of ethylamine, i s much
smaller than that of 1,3-diaminopropane.
These data are summa-
rized in Table 1 1 .
TABLE 1 1
RATE OF DIAMINE COMPLEXES AND THEIR STABILITY CONSTANTS
kKjXlO
5
_2
Log K
Q 1
Log K
1 2
Log K
mole
1*.
Q 2
sec"*
1
CH CH NH
3.37
3.93
7.30
1
H NCH CH NH
4.62
2.92
7.54
0.3
H NCH CH CH NH
5.77
1.42 )
3
2
2
1)
2
2
2
2
2
2
2
2
2
Results from the present study.
from Appendix I.
1
0.005
7.19 )
1
Other s t a b i l i t y constants are
It i s seen that K^ and kK^ both of which should reflect
2
(inversely) the chelating tendency of the mono-complex indeed
follow closely parallel trends.
1,4-diaminobutane which also has an abnormally large KQ^
value, similar to that for 1,3-diaminopropane, was expected
to show the same type of behaviour as the latter.
However, the
rate in this case was almost independent of the amine concentration and the kinetics were similar to those for ethylamine,
although the actual overall rate of reaction ( k
) *
w
e x p
8
almost
64
the same as for 1,3-diaminopropane (cf. Table 10).
The s i g n i f i -
cance of this behaviour i s not understood.
GENERAL DISCUSSION
A common feature of the systems examined i n the course of
this study i s that i n every case CO apparently reacts with a
species of the composition L-Ag-OH. The i n i t i a l reaction i n
each case can be represented as
L-Ag-OH + CO
> L-Ag-COOH
k
(ii)
The rate constant, k, of this process could not be measured
directly, but the data yielded values of kK^. In some cases,
the back reaction of ( i i ) was sufficiently fast to compete with
the subsequent reaction of the intermediate, L-Ag-COOH, with
another silver ion.
,
1
L-Ag-OH + CO „
k
-l
L-Ag-COOH
(ii )
1
A l l the values of kK^ (or k^K^) for various amine complexes
investigated i n this study are summarized i n Table 12, together
with the values of K^ K ^ and kK^K^K^.
s
d
The value of the latter
is identical to %k p i n the case of ethylamine-type "standard"
fiX
amine complexes.
For a l l the mono-dentate amines i t i s seen that
*
kK^ i s substantially independent of the nature of the amine.
* Triethylamine i s the only case where the deviation of the value of kK^ from the "standard" value of 1x10^ appears to l i e outside experimental error. For this amine the values of
and
both of which were used for the determination of kKjj were less
precise than those for the other amines; only one significant f i gure was available from the literatures. Furthermore, i n this system i t was necessary to use high (LH+} in most of the experiments
so that ionic strength effects which have not been taken into
account may be important.
65
TABLE 12
SUMMARY OF KINETIC AND RELATED THERMODYNAMIC DATA
xlO
sec"
xlO
mole'l"
4
2
Amine
1
xlO
mole'l"
xlO
xlO"
mole °1"2 mole°°2.l2.
sec"
8
4
1
1
5
2
1
0.03
1.2
0.18
1.43
2.9
5.2
0.78
1.2
6.5
7.6
1.0
H0C H.NH
2 4 2
0.14
2.8
0.55
1.5
1.0
(C H ) NH
4.3
5.0
9.1
(HOC H ) NH
0.16
NH
3
CH NH
3
2
C H NH
2
5
2
o
2
o
5
2
2
4
2
(C H ) N
2
5
120
3
(HOC H ) N
2
4
0.36
3
16
0.10
80
5.9
460
0.5
12
1.5
H N(CH ) NH
0.4
380
4.4
2
2
2
2
3
2
2
15
46
1.6
470
3.6
0.0079
H NCH CH NH
2
0.22
18
1700
1.4
0.9
1.0
1.0
2.5
1
0.30
0.0025
The constancy of kK^ over a 2,000 fold variation of K^K^ (and
hence of kK^K^K^) i s striking.
The constant kK^ may be identified with the rate constant
of the alternative and kinetically equivalent representation of
( i i ) , i.e., with the rate constant of the termolecular process
(ii"),
LAg
,
+ CO + OH"
kK
= ^ L-Ag-COOH
h
(ii")
It seems likely that K^ also i s insensitive to the nature of
66
L, and indeed probably does not differ greatly from the hydrolysis
constant of the free Ag+ ion whose value i s about 2x102 mole 1~1
(11). Using this value for K^, k i s estimated to be about
o - l - l
5x10* mole "l.sec .
This participation of hydroxide ion in the reaction
(base catalysis) accounts for the low reactivity of Ag
CO i n acidic media.
+
toward
The insensitiyeness of the reactivity of
L-Ag-OU to the nature of L, suggests that the amine molecule i n
L-Ag-OH i s acting only to solubilize AgOH and prevent precipitation of silver oxide.
The enhancement of reactivity i n these
amine-buffered systems would appear to be due mainly to the
high pH, rather than to specific complexing effects.
On the other hand, the rate of the back reaction of step
( i i ) and the rate and nature of the second step of the reaction
do appear to vary with the nature of amine, L. This i s shown
particularly by a comparison of the ammonia and triethylamine
complexes.
*
The apparent enthalpy and entropy of activation, A H
exp
and A
for k ^ ^ « kK^K^K^^ for "standard" amine complexes
listed i n Table 6 correspond to
A H
exp "
A s
e x p - ^S* + A S + A S + A
A
H
*
+
A
H
h
+
h
where AH
and AS
A
H
b
+
A
b
H
di
(3-46)
S d i
(3-47)
correspond to the enthalpy and entropy of
activation of the bimolecular process ( i i ) and the other terms
with subscripts h, b and d^ correspond to enthalpy and entropy
67
changes of the following equilibrium processes.
LAgOH - A H ^ - A $
h : LAg + OH" v
+
b sL + H0
d
±
: AgL
LH + OH" - AH^ - A S
v
2
(3-12')
h
(3-8")
+
b
LAg + L - A H ^ - A S ^
+
(3-7 )
+
2
B
Few thermodynamic data relating to these processes are available.
Only i n the case of ethylamine, has i t been possible to obtain
f a i r l y reliable values of A H ^ , A H , A S ^ and A S
using
1
1
available data (11, 12). The values of A H ^ and A H ^ thus
d
d9
d
obtained for ethylamine are 0.7 and 6.4 Kcal. mole" and A S ^
1
and A S ^ are -12.2 and 3.6 e.u. respectively.
A H^ and A S ^ as well as that of K
h
The values of
can be approximated
to that
for the free Ag ion as pointed out previously, which are e s t i +
mated to be about -2 Kcal. mole" and 4 e.u., respectively.
1
Hence, the enthalpy and entropy of activation of the bimolecular
process ( i i ) are estimated to be A H * ~ 9Kcal. mole" and
1
if
AS
n
<~ -15 e.u. (the bimolecular rate constant being k ~ 5x10
mole" -I"sec" at 2 5 ° C ) .
1
1
Although reliable values for AIL^,
A S d _ , etc. for other amines are not available, i t i s expected
1
•k
that A H and A S
&
for these systems w i l l not differ greatly
from those for ethylamine.
Peters and McAndrew (5) have recently reported the following
kinetics for the reaction of aqueous silver acetate with CO i n
acetate-buffered acidic media at 90°C. and high CO pressure.
68
+ k K (A )(A^OAc)(CO )
+
fi
3
[Kf-)
c
,
(3-48)
and interpreted these in terms of the following mechanism,
AgOAc + CO
=^->AgCOOAc
AgCOOAc 4- Ag(I)
Ag 4- CO 4- H 0
+
> Products
K
2
(slow)
c
(rapid)
AgCOOH 4- R"*~
AgCOOH 4- A g — 2 > 2Ag 4- C 0 4- H
+
(vi)
k
(rapid)
+
2
AgCOOH 4- AgOAc
(vii)
(slow)
2Ag 4- C 0 4- HOAc
2
(viii)
(ix)
(slow)
(x)
In this case there appears to be a contribution to the reaction,
not only from AgOH but also from AgOAc. The process ( v i i i )
which corresponds to ( i i ) , ( i i ) or ( i i " ) in the case of amine
8
complexes, i s faster than the subsequent steps and thus corresponds to a pre-equilibrium.
Hence, i t i s not kinetically dis-
tinguishable whether this i s a base-catalyzed process as in the
case of amine complexes or rather an acid-inhibited process;
in other wordsi whether the silver species which i s directly
reactive toward CO molecule i s a hydrolyzed species, AgOH, or
an unhydrolyzed ion Ag . In the case of amine-buffered solution
L-Ag-OH was seen to be the only reactive species.
I t i s of great
interest, therefore, to see i f the contribution of unhydrolyzed
Ag
+
i s detectable in the acetate-buffered solution.
It can easily be seen that rate constants and equilibrium
constants of these two processes, base-catalyzed ( i i " ) and acidinhibited ( v i i i ) are related as follows.
69
k,
k
K
FT"
=
(3-49)
Vw
-l
Suppose the base-catalyzed process ( i i " ) i s the only process
contributing to equilibrium ( v i i i ) , then the kinetics observed
by Peters and McAndrew requires that
ko (Ag"*1
l
«1
-l
and hence
2
(3-50)
J
k
k
l
K
f
a
k
^
8
^ «i k ^
(3-51)
From the data of Peters and McAndrew the value of k^K k2(Ag J/k ^
+
n
=
+
6
—2
—1
(= k^^fAg J/K^) can be estimated to be about 10 mole »l«sec
at 90°C. On the other hand the insensitiveness of the kK^ (rate
constant of the base-catalyzed process ( i i " ) ) for amine complex
to the nature of L over a 1,000-fold variation i n i t s basicity
suggests that the reactivity of uncomplexed Ag for the same
+
type of process should also be close to that of amine-complexed
species.
*
Using the value of kK^ at 25 C. and the previously
determined temperature coefficient, the value of kK^ at 90°C.
5 - 2
can be estimated to be 9x10 mole
not i n accord with (3-51).
2 - 1
»1 "sec . These results are
This suggests that i n the case of
* The value of
and A
are not reliable so that kK^ i s
to be used for comparison. The value for A H* +
ethylamine has been estimated to be 7.2 Kcal. mole" . This temperature coefficient being small and f a i r l y reliable, the resulting
kKh for 90°C. i s also f a i r l y reliable.
1
70
acetate-buffered solution there may indeed be a reaction path
involving direct reaction between unhydrolyzed Ag
represented by ( v i i i ) .
+
and CO
However, the rate constant of this
process i s not obtainable from these data.
This possibility that the unhydrolyzed Ag
+
ion, may also
be reactive toward CO i s not altogether unexpected, since,
as mentioned earlier, i t has already been found that in acidic
2+
solutions the reduction of Hg
by CO proceeds through reaction
with the unhydrolyzed ion (3). That such a path (i.e., OH independent) i s not observed in the case of the silver amine
complex may simply be due to the high pH of the solutions,
resulting in enhancement of the OH -dependent path.
Another interesting result of Peters and McAndrew's work
is the evidence suggesting that AgOAc i s also reactive toward
CO, presumably through an intermediate complex analogous to
L-Ag-COOH, i.e., Ag-COOAc. In this reaction the f i r s t step
(vi) i s rate-determining with a rate constant at 90°C. e s t i mated to be
-2
-1
-1
k^ = 3.6x10 mole 'l-sec .
This compares with a value of 2.5x103 mole -1^ l ' s e c -1 estimated
for the corresponding rate constant of L-Ag-OH toward CO, i.e.,
L-Ag-OH appears to be about 10"* times as reactive toward CO
as AgOAc. The activation energy of the process (vi) has been
determined to be 15 Kcal. mole" while that for the process ( i i )
-1
is 9 Kcal. mole
1
71
REFERENCES
1.
Halpern, J., Advances in Catalysis. XI, 301 (1959)
2.
Halpern, J . and Taylor, S. M., Disc. Faraday S o c , 29,
174 (1960).
3.
Harkness, A. C. and Halpern, J., J . Am. Chem. S o c , 83,
1258 (1961).
4. Bauch, G., Pawlek, F. and Plieth, K., Z. Erzbergbau und
Metallhutenwessen, XI, 11 (1958).
5.
McAndrew, R. T. and Peters, E., X l l l t h International
Congress of Pure and Applied Chemistry, Montreal,
Canada, August, 1961 and unpublished results.
6. Halpern, J . and Kettle, S. F. A., Chem. Ind., 668 (1961).
7.
Just, G. and Kauko, Y., Z. phisik. Chem., 82, 71 (1913).
8. Lattey,' J. Am. Chem. S o c , 1959, 29 (1907).
9. Von Georg-Maria Schwab, et a l . , Z. anorg. Chem.,, 252,
205 (1944).
10.
Seidel, "Solubilities of Inorganic and Metal Organic
Compounds" 3rd Edition, Vol. 1, p. 217 (1940).
11.
"Stability Constants of Complex Salts", special publication
of the Chemical Society.
12.
"Selected Values of Chemical Thermodynamic Properties",
U.S. Bureau of Standards circulation.
72
APPENDIX I
SELECTED THERMODYNAMIC PROPERTIES OF AMINES
AND SILVER-AMINE COMPLEXES
*2
Stability Constants*
*1
pKa
Amine
LogK-^ LogK
NH
*6
9.25
(9.09)
CH-NH-
LogK
12
*4
3
l
xlO
d
xlO
2
4
4
1.2
3.31
(3.24)
3.92
(3.81)
7.23
(7.05)
10.72
3.15
3.53
6.68
5.2
2.9
10.81
3.37
3.93
7.30
6.5
1.2
(C H ) NH
10.96
3.06
3.30
6.36
9.1
5.0
(C H ) N
10.77
2.6
2.1
4.76
5.9
HOC H NH
9.74
3.13
3.55
6.68
0.55
(HOC H ) NH
9.00
2.69
2.79
5.48
0.10
(HOC H ) N
7.90
2.30
1.34
3.64
7.9xl0"
H NCH CH NH
2
10.18
4.62
2.92
7.54
1.5
12
H N(CH ) NH
2
10.64
5.77
1.42*
4.4
380
H N(CH ) NH
2
10.82
5.9
3
C H NH
2
5
2
5
2
2
5
3
2
4
2
2
2
2
4
2
2
2
4
2
2
3
2
2
2
3
4
*1
(L) (H*"] / (LH+) - K
*3
(AgL )/(Ag )(L) = K
*4
(AgL^(L)/(AgL tl - K
+
5
0.174
(0.182)
7.19*
5
*2
+
2
2.8
16
460
3
(LH+nOH-VCL) = K
(AgL +J/(Ag L)CL3 - K ,
2
Q 1 >
80
6.6
a
+
(1.5)
12
K
b
- K^-K^
Q2
d i
These data are from "Stability Constants of Complex Salts", special
publication of the Chemical Society, corrected, where necessary, to
25°C. using the known temperature coefficient for Ag(NH C H5)2"*"
(d log K / d t - d log K /dt - -0.016
2
m
19
2
73
* 5
Estimated from the result of the present work (cf, p. 62).
* 6
Values i n parentheses are for 30°C,