Patrick G. McCormick
Marquette University
Milwaukee, Wisconsin 53233
Reaction Rate of Ethyl Acetate
Hydrolysis by Osdlometry
High frequency conductance, or 0scillometry, is usually introduced to the student in a
physical chemistry course, or one in instrumental
analysis. Especially in the latter case, not only is
the theory presented, but also the practical situations
in which the technique is applied to advantage. Among
the most important advantages of the technique is
complete avoidance of solution-electrode contact.
Another is the sensitivity to small changes in the sample
conductance and the speed with which those changes
can be detected.
Unfortunately, aU too often, upon entering the laboratory and finally confronting the instrument, the
student finds that the experiments he is to perform do
not take advantage of these features. Indeed, one of
the most common experiments performed with the
oscillometer is a titration, not even illustrating these
features. To be sure, the determination of loading
curves does illustrate pertinent theory, and measurement of dielectric constants and small concentrations
are typical applications of oscillometry. To the student,
however, they offer little challenge, and provide
little data.
A more pertinent experiment is the determination of
the second-order rate constant in the alkaline hydrolysis
of ethyl acetate. Such an experiment not only takes
full advantage of the unique features mentioned
previously, but also provides values which lend themselves to analysis by statistical measures. The experiment demonstrates vividly an important instrumental
method, offers a good introduction to automatic
recording of results, and provides data of real significance from which much additional information can
he extracted.
The a ~s~~ l i c a t i oofn hieh freauencv oscillators to the
measurement of rate constants has been described
previously (1-S), and their suitability for such determinations has been fully established. These workers
made use of rather elaborate temperature-regulating
equipment which pumped the sample solution (I)
or the thermoregulating liquid (2) through a cell
placed within the coil of a tuned oscillator circuit.
The presence of hydroxyl ions in the solution caused a
change in the capacitance of the solution and a corresponding imbalance was indicated in the oscillator
circuit. As hydroxyl ions were consumed in the
reaction, the circuit indicated a reduction in the imbalance condition. In effect, then, the oscillator
monitored instantaneously the concentration of hydroxyl ions in solution. This indication was of a
nature which could easily be displayed on a strip-chart
recorder providing a record of time versus a function
which could he related to concentration. Mathe~~L
-
558 / Journal of Chemical Education
matical analysis of the cuwe provided the desired
rate constant. The entire experiment took 20 minutes
or less.
Such an experiment seems ideally suited to the
student laboratory. The primary drawback is the
elaborate nature of the temperature-regulating and
circulating system. Nonetheless, an attempt was
made to adapt the method for use with the Sargent
Model V Chemical oscillometer, an instrument found
in many laboratories. During initial studies two
important discoveries were made. First, if all solutions, the instrument, and the cell are a t room tempera-.
ture to begin with, and one uses a large volume of
sample (ethyl acetate) solution, the addition of a
small volume of sufficiently dilute NaOH will not
cause any significant temperature change during the
relatively short time required for the reaction to take
place. It would appear that the cell plus sample
solution provide sufficient capacity for temperature
compensation under these conditions and that no more
elaborate regulation is necessary. I t should be noted,
of course, that one does not have the ability to choose
a particular temperature, being forced to work a t
ambient. However, one need not be concerned about
significant deviations from that. temperature during
the course of a single experiment.
Second, if the NaOH can be added and efficiently
stirred into the solution in the cellwithin the risetime
of the recorder (1-2 sec in this case), extremely wellshaped decay curves can be obtained which provide
data amenable to direct calculation of the rate constant.
With careful technique, the experiment gave theoretical results at two temperatures for which literature
values were available. In the hands of students,
excellent results were achieved, and in cases where
results deviated from theoretical values by more than
could he accounted for by experimental error, the
evaluation of probable sources of error provided
further insight into the experiment and the technique.
Literature values for the rate constant were available
for temperatures of 25' and 30°C. An Arrhenius
plot of log k versus 1 / T was prepared from these
two values and values a t other temperatures near
this range were obtained for comparison with experimental results. The agreement is quite good. (See
Table 2 below.)
Ethyl acetate was chosen for study in developing
this experiment, but other esters can undoubtedly
be used as demonstrated by Flom and Elving (2).
Conceivably, an unknown ester could be given and
its identity deduced from a calculation of the rate of
its alkaline hydrolysis.
Experimental
Procedure
Apparatus
A stock solution of 0.10 F NaOH was prepared and standardized
against potassium acid phthalate in the nsttsl manner. A stock
solution of ethyl acetate approximately 0.05 F was prepared by
diluting 10 ml of reagent grade material to 2 liters with distilled
water. The exact concentration of this solution was determined
by completely hydrolyzing a known quantity in a measured
excess of NaOH, hack-titrating the excess with standard acid.
This determination, due to volatility of ethyl acetate, should be
carried out every one or two days, since the exact value is needed
in the cdculrttions. These concentrations, though not those re*
ommended by Flom and Elving, were found to give the h a t
curves, and resulted in the best rate constant values
Prior to performing the experiment, a11 solutions and equipment
were allowed to reach temperature equilibrium for several hours
in the laboratory. The oseillometer was turned on a t least one
hour before use to reduce the likelihood of drift during a run.
Just before starting the run, the oscillometer meter was set to
mid-scale (zero current) using the center adjust and zero set
controls. When exactly balanced, the recorder pen was set a t
zero using the recorder zero control. The center adjust control
on the oscillometer was then used to move the meter needle to
ahout 16 to 17 pA to the left of zero. If the recorder did not
deflect upscale, the recorder-input leads were reversed. (Deflections on the meter during a run were always to the left of
zero.) The range attenuator on the recorder was used to bring
the pen to the top of the scale. This would allow the values ohtained during a run to use as much a9 possible of the chart.
The recorder chart meed was set to Medium I1
. in./min).
. .
E:xartly 100.0 ml of the ethyl acetate solution wa-5 pipetted
into the cell, ~ n rhe
d temperature was read with a thermometer.
The center ndjust and Zen) set mntrola on the oscillometer were
used to return the meter needle to zero. (Note that throughout
the entire experiment none of the variable capacitors are used,
nor is the internal standard.) If the recorder pen did not return
exactly to zero, it was adjusted accordingly.
The svrinee was filled with exsctlv 1 ml of base. no air bubbles
heing t;ap@d. The recorder was"started seveisl subdivisions
before a major grid line on the chart paper, and the syringe and
stirrer were poised over the cell, one in each hand. Just as the
pen reached the bare of a major uharr division, rhc contents of the
yringe were rapidly ejected direvtlg into the solution in the cell,
and the stirrer was nlunmd into the cell. to ahout ' 1 , of its denth.
One ouick un-anddown motion was d e . and the stirrer \as
removed from the cell. If done quirkly, the pen had nor renrhcd
the peak of its travel irp the chart lwfure the stirrer was removed
from the cell. A looj~filtingrap (an aluminum foil moisrure
cup serves very well) was carefully placed on top of the cell to
reduce volatility losses of ethyl acetate during the experiment.
The system was allowed to remain undisturbed for 18 to 20 ruin,
durine which time the recorder nroduced a. trace similar to
that bhown ill Figure 2. Thp rene;ion aaq considered rmnplete
when no change was observed in the recorded vnlue for 2 to 3
min.
The instrument used was the Sargent Oscillameter, Model V,
equipped with a large cell and holder. The instrument was connected to a Sargent SR recorder through the recorder jack on the
back of the ascillometer. A shielded cable was constructed which
connected to pins 1 and 3 of the female plug on the instrument,
the ground wire connecting to the instrument chassis via, the plug
casing. The recorder was used without a range plug, in which
configuration it had a 125-mV range. The oscillometer output
was still too large for accommodation by this range, so a 1:200
voltaee divider was constructed from orecision l-Mee and 5-
scale on the recorder by reducing the range by means of the range
~
ran be arr~mmodaredhy
attcnunrion vontrol. L R I ~ CrGrentq
using larger voltage dividers, I:&O0 heiug sl~ltir.ent to display
thr rntlre ourput of the o:(.illom~ter cimtir. It .h011111 tm mted
thnt rhic nrmngemntt COIIIIPCIS the reronlrr to thp o~rillomerrr
i r i a d n j u t t I n t . Thus, the revorder will follow
accurately rurrrnt vdurs whirh are olT+.nle on the mrter.
As noted above, two requirements were necessary to ensure
well-formed curves: rapid, accurate base addition, and rapid
stirring. Flom and Elving added base with s. pipet, starting the
recorder a t half-addition. In the present work, a more rapid
method was sought which offered at least the accuracy of pipets.
It was discovered that disposable plastic Tuberculin syringes'
without needles, worked extremely well. The syringe has a, one
ml capacity, and is not affected by the NaOH solution. Without
a needle attached, it om he emptied very rapidly. The graduations allow the plunger to he set quite reproducibly. To test
the accuracy of such a method, a randomly selected syringe was
calibrated for delivery using water and weighing the ejected
amount. Ordinary care was exercised, hut no special precautions or techniques were employed. As shown in Table 1, the
syringe showed an accuracy and precision superior to what might
be expected of a Class A l-ml pipet.
Table 1.
Calibration of Plastic I-ml Syringe
Trial
Weight dispensed (g)
Temperature = 2:3.0eC; density of water - 0.00757 e, ml.
averare weight drlivered = 1.00107 g; average volnme delivered
= l .(kl34 ml; s t d drv. = 0.0001; ml.
The design of the cell, heing of annular crass-section, makes stirring difficult. Glass stirring rods could damage the cell, and magnetic stirring would not he efficient, or rapid. A plastic stirring
rod was used with some success, hut, especially in the hands of
students, often proved inefficient as indicsted by poorly formed
curves. A plastic stirring apparatus was fabricated from Teflon
and Nylon as shown in Figure 1. The Teflon disk was a/* in.
thick, though this dimension is not critical. The Nylon handle
was made from '/&in. diameter stock, cut down to '/ain. at one
end to fit into one of the small holes as shown. The handle was
7 in. long. This stirrer fits into the annular cell, and two rapid
updown motions are sufficient to affect complete stirring of the
cell contents. This motion a n he produced, with very little
practice, in far less than the time required for the recorder pen
to resch its peak value.
'Available from American Hospital Supply, Evttnston, Ill.,
or from Pharmrtsesl Laboratories, Glendde, Calif.
Figure 1.
Dotoilr d plarticstirring apparatus.
Volume 48, Number 8, August 1971
Table 2.
2
TIME (MINI
Figure 2. Typieol recorder tracing during a wn 10 = 0.0482 F, b =
0.000986 F, tamp = 25'C).
Calculation
The data treatment is fully developed by Flom
and Elving (%), and will only be outlined here. The
curve is extrapolated to give values at time zero
(6) and infinite time (i,). The initial concentration
of base is proportional to lie - i-1, and the decrease
after time t is proportional to (io - it)/(& - i-),
which expression is written as a for convenience.
It is shown by Flom and Elving (%) that substitution
of this expression into the second-order equation
results in the following expression
where a and b are the concentrations of ethyl acetate
and sodium hydroxide, respectively. This expression
originally derives from MacImes (4) and was utilized
by Jensen, Watson, and Beckham (I) for similar
data. A plot of t versus the log term of the equation
should yield a straight line, the slope of which is
2.303/k(a - b) for a second-order reaction. Time
and current (arbitrary recorder units) data are read
from the chart, and the log term calculated for each
current value. These data are plotted and the slope
of the straight lme is used in determining the rate
constant. Since this is a very time-consuming task,
a short FORTRAN IV computer program was written
to take time and current data directly, and calculate
a least squares slope of the plot mentioned above.
From this value, the computer evaluates the rate
constant. (A copy of the program will be sent on
request.) It should be mentioned that lmearity in
the plot only extends for 13 to 15 min of data, so
points beyond that value should he checked for fit,
lest the slope be unduly influenced by them.
Results
As mentioned previously, values for the rate constant
for ethyl acetate hydrolysis a t 25°C and 30°C are
given in the literature. Theoretical values a t other
temperatures can be deduced from a plot of log k
versus 1/T constructed from these two known values.
In Table 2 are shown some typical results obtained
both by the author and by senior undergraduate and
first-year graduate students. Comparison is made
with theoretical values determined as described above.
It can be seen that values correspond quite well
with theory. It is expected that other esters would
give similarly satisfactory results. Flom and Elving
560 / lournal of Chemical Education
Rate Constant Values for Alkaline Hydrolysis
of Ethyl Acetate
Ternperature ('C)
This Work
(l/rnole-sec)
Theoretical (l/mole-sec)
24
25
28
30
32
0.104
0.110
0.134
0.145
0.158
0.104
0.110 (8, 8 )
0.132
0.145 (1, 9)
0.1.55
(9) give values for the rate constant of hydrolysis of
methyl and ethyl chloroacetate a t both 20°C and 30°C,
so these compounds might also be used as described
in this paper. The rates for these hydrolyses are
much faster than for ethyl acetate, and the experiment
would he expected to he proportionally more challenging.
Conclusion
As described here, the high frequency conductance
method can he used to determine a t ypical secondorder rate constant with very satisfactory results
by any careful student. The experiment offers many
significant improvements over other typical oscillometry student experiments (5-7) including (1) a clear
demonstration of some unique advantages offered
by the technique itself, (2) the opportunity for the
student to obtain very close to theoretical results
in a reasonable length of time, (3) the opportunity
to compare classes of similar compounds such as
acetates and chloroacetates, or ethyl acetate and
ethyl proprionate, and (4) the possibility of determining an unknown from a study of its kinetic behavior.
Additional features of the experiment include an
opportunity to use a computer as a computational
tool to greatly reduce the time required for calculation.
The program used in this calculation is not a t all
complicated and anyone with any FORTRAN skills
could easily write his own. Finally, by making
measurements of the rate constant a t two temperatures, additional thermodynamic information can be
deduced such as the energy of activation for the
reaction. The experiment is quite versatile, and can
be presented in several ways to serve the best interests
of a variety of course objectives.
Acknowledgment
The author wishes to express his appreciation of the
time and helpful suggestions offered by the students
who ran this experiment. Special thanks are due
P. J. Lamothe.
Literature Cited
(1) Jassew. F. W.. WATSON,G. M.. AND BEOIEAM,
d.
-1770
..- (10S1>
- -- .,.
B., ~ n d Chcm.,
.
23,
\
(2) Fmu, D. G., AND ELVINO,
P. d., A d . Cham., 25, 541 (1953).
(31 EWINQ,
P. J.. Awn LArsrrz, I., J . Amcr. Chcm. Soc., 77, 3217 (1955).
(4) MAOINNEB,D. A.. "The Principles of Electroohemistry;' Reinhdd
Publishiw Corp., N e w York, 1939, pp. 378--80.
(5) MELOAN,
C. E.. AND KIBER,R.
"Probleme and Experiments in Inatrumental Analysis." Merril Books. Inc.. Columbus, Ohio, 1963, pp.
W..
-.
,A""G G S
(6) W r m m o . H. H.,MERRIPT,L. L.. AND DEAN,J. A,, "Instrumental
Method. of Analysis," 4th ed.. D. Van Nastrand Co.. Prinoeton.
N. J . . 1965, pp. 741-2.
(7) D m * a ~ r ,P., "Instrumental Analysis," The Maomillan Co., New
York, 1957, p. 355.
, G., Tmna. Famday Soe., 40, 352 (1944).
(81 D n e r a ~ K.
(9) Rmcnss, L. T., Ann.. 238, 276 (1887).