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Application of Catalyzed Reactions in Analytical Chemistry.

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Application of Catalyzed Reactions in Analytical Chemistry
New analytical
By Herbert Weisz[*]
I
I
Catalyzed reactions are becoming increasingly important in analytical chemistry, and especially
in trace analysis. The acceleration of a particular reaction (indicator reaction) promoted by
the catalyst provides information about the catalyst concentration. Substances which lower
or raise the activity of a catalyst (inhibitors or activators, respectively) can likewise be determined
in this way. The course of a reaction is usually monitored by physical measuring techniques.
Various methods of conducting these kinetic-catalytic studies are discussed in this article.-Use
of catalyzed reactions for end point detection in volumetric analysis is also feasible.
1. Introduction
Whenever a definite relationship exists between the extent
of a physical phenomenon and the amount of substance giving
rise to this phenomenon, measurement of this property can
be employed for quantitative determination of that substance.
All measurable properties can be utilized for this purpose.
During the measurement, the system must be in a state of
sufficient stability, i. e. in thermodynamic equilibrium.
On the other hand, precisely the change of a property
of a system with time may be monitored, i.e. time may be
employed as an additional measured quantity.
Kineticanalytical techniques of this kind have found increasing application in recent years. Detailed accounts are to be
found in the monographs by Frost and Pearson"] and Mark
and Rechnid'], and especially on practical analytical aspects
in the book by Y~tsirnirskii[~~;
some more recent surveys should
also be m e n t i ~ n e d [ ~ - ~ ] .
An important group of kinetic analytical methods is based
on the use of catalyzed reactions.
These techniques attract particular interest in analytical
chemistry, doubtless because they provide a highly sensitive
means of detection or quantitative determination of the catalyst itself.
Catalysts are substances which increase the rate of a chemical reaction. The requisite lowering of the activation energy
proceeds in such a way that the catalyst is regenerated and
thus maintains a constant concentration in the system; hence
even small amounts of catalyst can effect high conversions,
thus accounting for the high sensitivity of catalytic-kinetic
methods.
Catalytic reactions have long held a firm place in qualitative
analysis. Mention need only be made of the extremely sensitive
iodine-azide reaction for the detection of ~ulfide''~;iodine
and sodium azide practically fail to react with each other,
reaction occurring only on addition of sulfide (or various
other compounds of divalent sulfur):
It is readily observed, both by the disappearance of the
brown color of the iodine and by the development of gas
bubbles (nitrogen).
[*] Prof. Dr. H. Weisz
Lehrstuhl fur Analytische Chemie
Chemisches Laboratorium der Universitat
Albertstrasse 21, 7500 Freiburg (Germany)
150
Analytical methods based on the use of catalyzed reactions
are of great importance in present-day trace analysis, primarily
owing to their frequently marked selectivity (especially in
enzymatic analytical methods) and their high sensitivity, comparable with those of activation analysis.
Let us consider the reaction of two substances accelerated
by a catalyst K
A+B
K
-+
X+Y
The concentration of K can be determined from the extent
of this acceleration. The procedure adopted entails monitoring
the rate of reaction between the substances A and B for
a change in concentration of one of the starting materials
(A or B) or of one of the reaction products (X or Y).
The reaction employed for determination (or detection) of
the catalyst is called an indicator reaction. That substance
whose change in concentration serves as a measure of the
progress of the reaction is termed an indicator substance.
It may be either a reactant (A or B, iodine in the case of
the iodine-azide reaction) or a product (X or Y, elemental
nitrogen in the case of the iodine-azide reaction).
The progress of the indicator reaction can therefore be
followed by monitoring the depletion of a reactant or the
accumulation of a product.
Assuming the indicator reaction to obey the following rate
law:
k=rate constant; Cq, C,=concentration of reactant A and B respectively
at a time r
then the reaction is of second order since the catalyst concentration of CKremains unchanged.
Now if one of the reactants, say B, is present in such
excess that its concentration can also be regarded as constant,
then we have a reaction of pseudo-first order:
Ci, C,O=concentration of A and B, respectively, at time t = O
h'= k C $
If it is possible to determine even very small amounts of
the indicator substance X by a sufficiently sensitive method
of measurement (a small depletion of A is much more difficult
Anyeti. Chem l n t . Ed. Enyl.
1 Vol. 15 (1976) No. 3
to determine) then determination of the reaction rate can
be restricted to the initial state of the reaction. This means
that during this brief period the concentration of reactant
A can also be viewed as constant and we have a reaction
of pseudo-zero order:
x
dc
dt
~
k'C,Ci=k'CK=constant
(k"= k ' C 9
(3)
We see that in the initial state the reaction rate is directly
proportional to the catalyst concentration C,. Integration
of this relation affords:
cx% k ' C,t
(4)
Equations(3)and (4)are the basis of the so-called differential
methods. Three basic methods are available for determining
the catalyst:
1) Tangent method: the change in concentration of X (CJ
is continuously observed and registered uersus time. The slope
of the resulting straight line, tancr, is a measure of the catalyst
concentration (C,). In the calibration graph the tancc values
are plotted against the corresponding catalyst concentrations.
2) Fixed time method: the amount of reaction product
X formed within a fixed, predetermined time interval from
the start of the indicator reaction is determined. In the calibration graph the Cx values are plotted against the corresponding
catalyst concentrations CK.
3) Fixed concentration or variable time method: the time
which elapses from the start of the reaction until Cx has
attained a certain predetermined value is measured. The reciprocal of this time plotted against the pertinent catalyst concentration affords the calibration graph.
If measurements have to be extended beyond the initial
state, i.e. if the concentration of the reactant A present in
lower concentration can no longer be regarded as constant
(after consumption of more than 10% of A), then integral
methods are used instead of differential ones. They are based
on the integrated form of the differential rate law [cf. eq.
@)I:
There is no longer a linear relation between Cx and t
[as in eq. (4)] but a logarithmic one instead. Nevertheless,
the three basic methods just described can still be employed.
By suitable choice of reaction conditions (concentration,
pH value, temperature, ionic strength) the order of reaction
can be modified in such a way that a simple, optimally linear,
relation exists between the catalyst concentration and the
reaction rate. Most of the indicator reactions used in analytical
practice are modified so that they approximately correspond
to the above conditions of a pseudo-first order reaction.
(Theoretical aspects receive detailed treatment in the monographs" 31.)
The most frequently used technique for determining the
catalyticactivity of a sample is the following: The two reactants
of the indicator reaction (one generally in great excess) are
brought together and the reaction initiated by admixture of
the catalyst to the determined. In some cases it is more convenient to first mix one reactant with the catalyst and then
-
A i i ( I w . Chrvir. I n t .
E d . E J I ~ I/ . I'd. 15 (1976) N o . 3
start the reaction with the aid of the second reactant; this
procedure is recommended above all when the contribution
of the uncatalyzed reaction cannot be neglected.
The course of reaction can be followed by chemical or
physical methods.
The chemical approach is practicable only in the fixed
time method: the reaction is stopped after a suitably chosen
time; this can be achieved by rapid cooling or a drastic pH
change or, most frequently, by addition of an inhibitor which
has a sufficient "braking" effect on the catalytic activity. Then
either the unconsumed component of the reactant serving
as indicator substance or the product formed since the start
of reaction can be determined chemically (generally by volumetric analysis).
Measurement by physical techniques can be applied to
all three methods-tangent method, fixed time method, fixed
concentration method. This approach entails measurement
of the change in some physical property of the system which
is proportional to the change in concentration of the indicator
substance concerned.
In most of the examples given in the literature, the reaction
rate is determined by photometry in the reaction vessel itself
which serves as a measuring cell in the spectrophotometer.
Other frequently used methods include highly sensitive
fluorescence measurements, potentiometry, biamperometry,
conductometry, and, especially recently, enthalpimetric monitoring of the reaction course by thermometric methods; the
last method is particularly versatile, thanks to the sensitive
thermistors nowadays available.
Since the reaction rate is extremely dependent on temperature (cf. Arrhenius equation) the reaction vessel must always
be well thermostated; the enthalpimetric method clearly represents an exception where, by way of contrast, adiabatic reaction
conditions must be employed (e.g. in Dewar vessels).
The results are usually evaluated by graphical extrapolation
from a calibration graph.
Now not only the catalysts themselves can be determined
by these kinetic methods but, indirectly, any substance which
may influence catalytic activity :inhibitors, activators, reactivators.
Inhibitors, which lower the activity of a catalyst (e,g. by
precipitation or complexation), are determined by adding the
inhibitor to a known amount of catalyst and establishing
the fall in the reaction rate by one of the above methods.
A fitting example is the determination of nanogram quantities of mercury by inhibition of the iodide-catalyzed ceriumarsenic
(see Section 2.1) (1 ng= IO-"g).
Reactivators (sometimes called pseudoactivators) can reactivate, i. e. liberate (e. g. by demasking), a catalyst present in
an inactive inhibited form (e.g. as a complex); measurement
of the activity of the reactivated catalyst affords a means
of analytically determining such substances: manganese([[)
catalyzes the oxidative decoloration of alizarin S by hydrogen
peroxide; truns-1,2-diaminocyclohexane-N,N,N',N'-tetraacetic acid inhibits the manganese catalyst by complexation. Now
since lead is able to displace manganese from this complexand thus regenerate its catalytic activity-microgram quantities of lead can be determined indirectly by this methodL8]
(1 pg=IO-hg).
Activators enhance the activity of a catalyst and therefore
become accessible to determination. For instance, arsenic(rrr)
enhances the catalytic action of osmium on the reaction
151
between bromate and iodide and has been determined in
this wayL9!
In principle, a reactant of a catalyzed reaction can also
be determined-albeit
with a predictably lower sensitivity.
This possibility is only employed in enzymatic analyses (“determination of substrate”).
The reviews
list numerous examples of such reactions, and a selection of recent developments is presented
in Table 1.
Table 1. Some examples of redox reactions.
__
______________
Substance
Reaction
Ref,
Vanadium in blood
p-Hydrazinobenzenesulfonicacid
+ClOj
o-Dianisidine + H,O,
1- + H 2 0 ,
Alizarin S H,O,
p-Phenetidine H L 0 2
o-Dianisidine + 10;
3-Naphthylamine NO;
As”’ + BrO :
Alizarin red S + H , O ,
[I21
____
2. Indicator Reactions
The reactions used for determining catalytic activity must
satisfy a number of conditions:
The extent of uncatalyzed reaction should be as low as
possible (ideally zero).
The course of reaction should be easy to monitor, wherever
possible by several methods (e.g . photometry, potentiometry,
thermometry).
The catalyzed reaction should not be too fast, but, for
practical reasons, not too slow either; if it is too fast (tl,,< 10 s)
the time required for thorough mixing of reactants can no
longer be neglected.
Hitherto redox reactions and enzyme reactions have been
employed most widely as indicator reactions. In addition,
limited use has also been made of homogeneous catalyzed
exchange reactions of complexes with mono- and multi-dentate
ligands.
Concerning the heterogeneous catalyzed electrode reactions
already described by Heyroosky“ (catalytic “waves” in polarography) the reader is referred to the literature.
2.1. Redox Reactions
The principal oxidizing agents used are: hydrogen peroxide,
peroxoborate, oxygen, oxoanions such as bromate, iodate,
nitrate, peroxosulfate, periodate; the reducing agents include
iodide, thiosulfate, arsenate(m), a series of organic dyes such
as alizarin, malachite green, azo dyes ( e . g . methyl orange),
other organic substances such as hydroquinone, a-naphthylamine, p-phenetidine, o-tolidine.
Catalysts that can be directly determined by these reactions
include: copper, iron, manganese, titanium, zirconium(iv),
thorium(rv), molybdenum(vr), tungsten(vr). Redox reactions
catalyzed by osmium and ruthenium have been studied particularly thoroughly.
Other redox systems have also been employed for kinetic
determination of catalysts, for example the cerium-arsenic
reaction proposed as long ago as 1934 by Sandell and Kolthoff[’ ‘1; it is catalyzed by iodide and by osmium:
2CeiV+As”’
yellow
1--f
0s
+
+
+
S ,O 3
Copper
~
+ Fe”’
~ 3 1
1141
181
~ 5 1
1161
“71
I181
1191
POI
An interesting group of redox reactions warranting mention
are the peroxide oxidation of luminol or lucigenin catalyzed
by various metals (copper, manganese, chromium, cobalt) during which chemiluminescence occurs. The light emitted can
easily be measured by photometry (possibly with the aid of
a photomultiplier). In this way 26 pg/ml of chromium(i1r) are
determinable by a flow technique[’’] (1 picogram = 10-’2g).
2.2. Ligand Exchange Reactiorrs
A fitting example is the silver ion-catalyzed exchange reaction between hexacyanoferrate(r1) und 2,2’-bipyridine
(bipy)’ 2 ] :
’
[Fe(CN),I4-
+ 3 bipy
Ag +
[Fe(bipy),]’+
+ 6CN-
In similar manner, traces of mercury in the air have been
determined by hexacyanoferrate and nitro~onaphthol[~~!
An exchange reaction involving complexes with multidentate ligands is the following process which has been called
a “coordination chain reaction”: copper-edta and and nickeltrien react together via ligand exchange to form copper-triene
and nickel-edta; the reaction rate is greatly accelerated by
small amounts of additional complexing agent ethylenediaminetetraacetic acid (edta) or N,N’-bis(2-aminoethyI)ethylenediamine (“triethylenetetramine”, trien). The course of reaction can be conveniently followed by
2.3. Enzyme Reactions
2Ceiii+AsV
colorless
This reaction has since found numerous applications, being
an especially favorable indicator reaction: it can be readily
monitored by photometry, potentiometry (change in CeIV:
Ce”’ ratio), biamperometry, and thermometry.
The iodine-azide reaction mentioned in Section 1 which
is catalyzed by divalent sulfur has also been frequently utilized.
Some metals such as copper and manganese catalyze the
decomposition of hydrogen peroxide; this reaction can be
followed particularly well by thermometry.
152
Chromium
Molybdenuminplants
Manganese
Iron in plants
Ruthenium in minerals
Osmium, ruthenium
Osmium
Cobalt in reactor cooling
water
The determination of enzyme activity is based on the same
principles as that of the catalytic activity of inorganic substances. Enzyme reactions are employed primarily in clinical
and biochemical laboratories.
Apart from enzyme activity, substrates, i. e. reactants, can
also be determined highly selectively e. g . glucose by means
of glucose oxidase:
Glucose + H,O
+ 0,
H2O2+2I- +2H+
glucose oxidase
+
Gluconic acid
+
I,+2H,O
Mo
+ H ’0
Angew. Chem. l n t . Ed. Engl. f Vol. 15 (1976) No. 3
This represents an example of the coupling of two catalyzed
reactions. The rate of iodine formation provides a measure
of the glucose concentration[2’’.
As in the analysis of other catalysts, any substance having
a positive or negative effect on the catalytic (here enzymatic)
activity, can be determined. Various metal ions inhibit
enzymes; for instance, pg quantities of silver or mercury can
be determined by inhibition of alcohol dehydrogenasel2‘?
Moreover, non-metals can also be determined by this method:
pig liver esterase is significantly inhibited by traces of fluoride(*’I. Particular significance may be attributed to the quantitative determination of pesticides (e.g. Lindane, Aldrin, DDT)
by way of their inhibitor action on enzymes such as lipase,
phosphatase, and above all cholinesterase’**!
Reactivation of inhibited enzymes can be utilized for quantitative determination of the reactivator; thus traces of cyanide
and sulfide are determined by measuring the reactivation of
invertase inhibited by r n e r c ~ r y ~ ~ ~ ~ .
O n removal of the metal ion from a metalloenzyme (e.g.
zinc from alkaline phosphatase by ethylenediaminetetraacetic
acid) there is a loss of enzyme activity; slight traces of the
Same metal (in this case zinc) reactivate the enzyme. This
provides a highly sensitive determination of the activator[30!
The extensive field of enzymatic analysis is considered at
great length in the books by berg me ye^-[^'] and Guibuu/t[281;
just a few indications of possible application will suffice in
this article.
Since all reactions proceed simultaneously under the same
conditions there is no need for thermostating; moreover,
neither time nor concentration have to be measured directly.
While this ‘‘simultaneous comparison method” probably
represents the nearest approach to the “tangent method”,
systems of the “Landolt type” provide an opportunity of carrying out processes which can be regarded as “fixed concentration methods”.
A slow reaction is connected with a fast one by the product
X of the former (slow) reaction. Hence it follows that X can
only be present in the free state once all the R has been
consumed, i. e. the system must be deficient in R.
A+B
X+R
+
--t
X + Y slow
E
fast
If the first reaction can be accelerated by a catalyst, then
the time elapsing between the start of reaction and the appearance of free X, the “induction period”, is a measure of the
catalyst concentration: the shorter the induction period the
greater is the catalyst concentration.
This “chronometric” method has been successfully applied
to the determination of a series of catalytically active metal
ions (Fe, Mo, Zr, V)IJ3! For example:
BrO;+5Br-+6H+
3 Br,
+ 3 ascorbic acid
V
pH-5
+
6Br-
3Br2+3H20
+ 6 H + + 3 dehydroascorbic
3. DeterminationsEmploying Closed and Open Systems
acid
A whole series of possible methods are available for the
quantitative interpretation of the catalyzed reaction for determining the catalyst. An initial distinction is made between
“closed” and “open” systems: closed systems are those in
which there is no external intervention during the reaction;
neither reactant nor product are added or removed once
the catalytic reaction has started.
In open systems, on the other hand, reactants are added
and/or products removed during the reaction.
This is a self-regenerating system, the reactant B (Br-)
of the first reaction being re-formed in the second reaction.
Only after all the ascorbic acid has been consumed does
bromine appear momentarily and can be detected, for example,
with a fluorescence indicator (trypan red)r341.
The next procedure we shall consider is a catalytic difference
method: two vessels contain equal amounts of a reaction
mixture A + B and the catalyst serving to accelerate the inhibited reaction, albeit at different concentrations (sample K ,
and standard K J ; thus the two systems differ only in the
concentration of catalyst
3.1. Closed Systems
The general procedure described in Section 1 illustrates
the use of a closed system. Let us now briefly consider three
further methods of this kind.
The visual procedure referred to as the “simultaneous comparison method” by Bognar (1963)[321is based on a simple
experimental set-up described as early as 1922 by Huhn and
Leimbach: equal volumes of a solution of reactant A and
increasing known amounts of catalyst K are pipeted into
a series of test tubes; the sample solution having an unknown
concentration of K is then placed in one of these test tubes.
Immediately, equal volumes of a solution of reactant B are
added simultaneously, thus starting the reaction in all cases.
Observation of changes occurring in the individual mixtures-e. g. coloration, decoloration-reveals which of the
standards corresponds most closely to the sample in its behavior. Under identical conditions, equal changes clearly also
mean corresponding catalyst concentrations. This method utilizes a “dynamic colorimetric standard scale”.
Angew. Chmni. l n r . Ed. Engl. f Vul. 15 ( 1 9 7 6 ) No. 3
X
sample:
A +- B -L X + Y
standard:
A + B 2 X+Y
x
Whatever property ofthe system is employed for monitoring
the reaction course-e. g. the extinction or the potential-the
difference between the two measured values will always initially
be zero. Subsequently, however, the differing catalyst concentrations in the two mixtures will lead to a steadily increasing
difference which rises to a maximum; finally, after reaction
has reached completion in the two systems (this may take
a considerable time) the difference is again zero.
This change in the difference with time will provide some
information about the ratio of the catalyst concentrations,
either from the height of the maximum, from the area under
the curve, or from the slope of the curve leading up to the
maximum; hence it should be possible to establish the
unknown concentration in the sample.
153
Now if A is supplied at the same rate at which it is consumed
in the reaction, it follows, because C4 is no longer a variable,
that:
0
50
100
150
200
t
[
S
l
250
300
350
L
Fig. 1. Variation or extinction difference with time.
Figure 1 shows a curve obtained in such an experiment,
namely the photometrically recorded difference in extinction
in the copper-catalyzed reaction
The intensely red complex of iron(1ri)with thiocyanate serves
as indicator of the course of reaction.
This reaction was utilized for the determination of copper
(1-25pg in 1Om1; standard always 5pg). It proved most
convenient to employ the maximum difference in extinction,
occurring within minutes, in evaluation of the results; once
this value has been recorded measurement can stop, there
being no need to record the entire curve. The desired copper
concentration can be extrapolated from a calibration curve.
Measurement can readily be performed in a Lange colorimeter‘*’. 35!
In such difference determinations the reaction course can
also be followed by potentiometric’zOl,thermometric, and conductometric methods[”].
3.2. Open Systems
Flow, “steady state”, and “stat” methods should be mentioned under this heading. They all have in common that
the course of reaction is subjected to external influence by
addition of reactants and/or removal of products.
In the “stat” methods[371a given state of the system is
arbitrarily imposed and maintained during the reaction by
addition of a reagent. The system itself controls this addition;
the rate of addition of reagent necessary for maintaining the
given state provides a measure of the reaction rate and hence
for the catalyst concentration to be determined. The factor
maintained, and thus serving the control function, is a physical
property of the system dependent upon the concentration
of a reactant, a product, or the ratio of reactant to product.
Let us again consider the general reaction
A+B
K
-t
This is therefore a reaction of “quasi”-zero order : the reaction rate remains constant over a relatively long period, i. e.
beyond theinitial state,and the rate of addition is proportional
to the catalyst concentration CK.
An example utilizing photometric control may serve as
an i l l ~ s t r a t i o nmanganese
~ ~ ~ ~ : catalyzes the oxidation of malachite green by periodate (at pH= 3-4), reaction being accompanied by decoloration:
10;+ malachitegreen
colorless product
A schematic diagram of the apparatus, termed “extinctiostat” (“absorptiostat”) is shown in Figure 2: the thermostated
r-l
Recorder
C
F
O
C
111(10.11
P
Fig. 2. Extinctiostat (absorptiostat). L, light source; F, monochromatic filter:
B, iris diaphragm; C, lhermostated cell; P. photocell; A, amplifier.
cell C contains the reaction mixture (malachite green + 10;)
with the Mn solution to be determined; L is the light source
(here 619nm). The signal from the photocell (P) is amplified
(A), registered as a potential by a mV meter and fed to the
control device. Here the measured value is compared with
a given set potential. This set potential corresponds to a
certain set extinction in the cell and thus logically to a definite
selected concentration of malachite green. The difference
between the set and the actual potential, and thus between
the set and actual concentration of malachite green, arising
in the course of the reaction, is corrected by a command
to the buret: the deficit at any time is compensated by addition
T
-
360-
u
X+Y
240-
and assume B to be present in considerable excess so that
its concentration can be regarded as constant; then, if the
reaction course is followed by measurement of A:
11100.31
d C s - kCKCA
dt
154
Mn
--+
I
2
3
L
5
6
I
8
Halacmie green SOlVtiDn I m l l d
9
10
Fig. 3. Curve plotted for 7.5 pgMn145 ml catalyzing the oxidation of malachite
green by periodate.
Angew. Chem. I n r .
Ed. Enql.
Vol. 15 ( 1 9 7 6 ) No. 3
of a corresponding amount of malachite green standard solution. The recorder registers the rate of addition (ml/min) of
standard solution; this provides a measure of the catalyst
chamber
concentration (Mn) to be determined. A curve of this kind
as registered by the recorder is shown in Figure 3. Evaluation
is based on the cotangent of c( (vol/time) taken from the
plot. A calibration graph is plotted with several cotangent
CL values for known manganese concentrations.
and methods employing a potenThe pH stat
tiostatiJOl,a b i a m p e r o ~ t a t ’ ~and
~ ] , a l u m i n ~ s t a tall
~ ~function
~l
- - - - _ _ Point of
measurement
in analogous manner; in the last-named case the chemilumines1
-@7
jJ
cence occurring on reaction of luminol with hydrogen peroxide
Fig. 4. Schematic representation of flow tube. A, H. reactants: K . catalyst.
ismaintained constant by controlled addition of H 2 0 > Metals
catalyzing this reaction, e. g. copper, can be determined by
this technique.
The
reactants A and B and the solution
of the catalyst
While all these “stat” methods involve addition of some
K to be determined flow separately into amixing chamber,
reagent controlled by a change of a physical phenomenon
whence they enter a flow tube. At a given point, some property
in the system itself, “steady-state’’ methods entail admixture
of the system (potential, optical density, etc.) is continuously
of a reactant at a constant rate. This technique was employed
monitored. The
distance between
the mixing chamber and
in a similar form by Pearson and Piette[431for determining
the point
of measurement corresponds to a certain fixed time,
kinetic data of noncatalyzed reactions. T o a system containing
provided of course that the flow rate is constant.
one of the two reactants (A) in great excess as well as the
The value measured provides a measure of the catalyst
catalyst (K) to be determined is added
the other reactant
concentration.
(B) at constant rate by means
of a motor-driven piston buret.
The flow technique is well suited to continuous
analysis,
The reaction A + B starts directly upon addition of the first
any change in catalyst concentration will clearly alter the
small portion of B, albeit very slowly owing to the initially
value measured.
very low concentration of reactant B in the system. This
Although this kind
of kinetic method has been known for
means, however,
that not even the minute amount of
B will
over five decades there are only few analytical applications.
be completely consumed.
used such an experimental
set-up for continuous
Since B is added continuously, its concentration anddetermination
hence
of molybdenum
(1-50 pg) with the aid of the
also the rate of reaction A + B will
increase; and
so will
reaction of peroxoborate with iodide.
the consumption of B.
Whereas steady-state, but different, concentrations of reacAfter some time a stage is reached at which precisely the
tants and products will be established at different positions
same amount of B is consumed per unit time as is added
of the flow
tube if the reaction conditions remain unchanged,
from the buret. A steady state of the system-with regard
that variant of the flow technique employing thorough back
to the concentration of B-is then attained and will persist
mixing (“capacity flow ~ e l l ” ) [is~characterized
~1
by attainment
for some time.
of an effectively uniform steady-state concentration of the
Whatever physical property, whatever phenomenon is
reactants and products throughout the entire reaction vessel.
chosen for measuring the steady-state concentration of B,
This method will now be illustrated by the manganese-catae.g. the extinction or the potential, it serves as a measure
lyzed oxidation-and hence decoloration-of alizarin S by
for the reaction rate and thus, under otherwise identical condihydrogen pero
tions, as a measure
for the concentration of catalyst to be
determined. As always,
a calibration
graph is used in evaluaK
tion; the reciprocal value of the steady state concentration
(or the corresponding measurement, e. g. extinction or potential) is plotted versus the catalyst concentration.
In this way it was possible, fortoexample,
determine
between 60 and 650ng
of iodide using the Sandell-Kolthoff
reaction between arsenic(1rr)and cerium(rv): the iodide
sample
is added to an excess of arsenic(lr1)
and an approximately
M solution of cerium(1v) added from a motor-driven
piston buret at a rate of about 10 pl/min to a reaction volume
of about 10ml. The change in optical density (at 420nm)
is recorded. After a few minutes a steady
state
is attained,
the extinction remaining unaltered[44!
In the methods just described one of the reactants is added
to the other,
the rate
of addition being either constant or
Fig. 5. Schematic representation of a flow cell. A, H, K. see Fig. 4.
regulated by the system itself.
It is, of course, also possible not
only to add a reagent
but
to simultaneously remove an equal volume of reaction
the reactants alizarin S and hydrogen peroxide flow
separately
This is practiced in flow methods. A flow apparatus
solution.
5) where they are immediately thorinto a flow cell (cf. Fig.
is schematically depicted in Figure
oughly mixed. After some time a steady state is established
4.
‘tt
Aiipw C h i . liii.
E d . Engl.
Vol. IS (1976) N o . 3
155
with regard to the concentration of all reactants and products.
The concentration of alizarin S is continuously monitored
by photometry and recorded. It affords a measure of the
amount of catalyst: the higher the concentration of manganese
the lower the concentration of alizarin S.
As with all other catalytic-kinetic methods, these flow
methods can also be used for determining inhibitors, activators,
reactivators.
Special flow techniques have also been developed for following very fast reactions (e.g. stopped flow, accelerated flow
methods)[*];however, they have only rarely been employed
for analytical purposes.
The flow technique is also well suited to automation. A
detailed report has recently been published on automation
and instrumentation of kinetic analytical procedures employing open and closed systems1471.
4. Catalytic End Point Indication in Volumetric Analysis
All the methods discussed so far require quantitative determination of catalytic activity in order to determine the catalyst
itself or, indirectly, an inhibitor, activator, or reactivator. On
use of catalyzed reactions for end point determination in
volumetric analysis, it suffices to establish qualitatively that
the catalyst is present at a concentration high enough to
detectably enhance the rate of an indicator reaction. During
the past few years catalytic reactions have been employed
for indication purposes in volumetric determination^[^! A
review of this comparatively young field has recently been
written by Hadjiio~nnou[~~l.
The volumetric solution (K) which catalyzes the reaction
between substances A and B is added from a buret. This
volumetric solution can be used for titration uersus a substance
(an inhibitor) which drastically impairs the activity of the
catalyst K, possibly by precipitation or complexation. The
first excess drop of solution from the buret (i.e. of catalyst
K) accelerates the indicator reaction between A and B; this
is exploited for end-point detection.
The method is illustrated by the precipitation titration of
silver with potassium iodide standard solution: to the silver
solution to be determined is added yellow cerium(iv) sulfate
solution and arsenic(ii1); titration with 0.1 N potassium iodide
is then performed with vigorous stirring( !). Iodide ion is precipitated as silver iodide. The first excess potassium iodide
catalyzes the cerium-arsenic indicator reaction and the yellow
cerium(1v) solution is immediately d e c ~ l o r i z e d ~ ~ ~ l .
Chloride, bromide, iodide, and thiocyanate can also be
determined in this way by back titration.
Such catalytic end-point determinations have the advantage
that the excess of titration reagent is not consumed in a
stoichiometric reaction with an indicator but merely acts as
a catalyst; a very small excess of volumetric reagent is able
to catalyze reactions of comparatively large amounts of the
reagent mixture serving as indicator.
An example of reducing catalytic activity (inhibition) as a
result of complex formation is provided by the titration of
ethylenediaminetetraacetic acid with a cobalt volumetric
solution; oxidation of Tiron by peroxoborate (gold coloration)
serves as indicator reactionr491.
The end point does not have to be determined visually;
any physical method permitting determination of changes
156
accompanying the occurrence of the indicator reaction in
the system can be utilized. Thus catalytic end-point determinations employing p o t e n t i ~ m e t r y ~biamperometry,
~~l,
photometry["], and
5 3 1 have all been described.
The last-mentioned procedure appears especially interesting
since it is not the slight evolution of heat during the actual
titration reaction (catalyst + inhibitor) which serves for endpoint recognition, but instead the much greater one accompanying the indicator reaction (A B) which occurs once the
end-point is passed.
+
Received: December 15. 1975 [A 100 IE]
German version: Angeb. Chem 88, 177 (1976)
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/ Vol. IS
( 1 9 7 6 ) No. 3
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