C. E. Hedrick University
of
Pennsylvania Philodelphia
I I 1
Formation of the Chromium-EDTA Complex A n undergraduate kinetics experiment
W h e n a solution of a chromium(111) salt is added to an excess of a solution of disodium EDTA [ethylene(dinitrilo)tetraacetate] at approximately pH 4, apparently no reaction occurs a t room temperature. After a short time, however, the light gray-green color of the chromium(II1) aquo ions changes to the deep purple color characteristic of the chromium(111)-EDTA complex. This reaction is used as the basis for an experiment in reaction kinetics for an undergraduate course in analytical chemistry. Hamm has studied the absorption spectra of the chromium(II1)-EDTA complex, and finds that the formation of the complex involves several slow steps.' The first slow step is first order with respect to chromium(II1) and inversely proportional to the hydrogen ion concentration. The slow rate of the reaction is not expected from the large value of the formation constant of the complex. The logarithm of the formation constant is 23, as reported by R i n g b ~ m . ~ An example of the absorption spectrum of the chromium(111)-EDTA complex, obtained with a Spectronic 20 spectrophotometer, is shown in Figure 1. The purple color actually measured represents a mixture of several intermediates. Hamm reports that at 545 mfi, the molar absorptivity of the compound HCr(EDTA)(H,O) is 200 at pH 5, while the molar absorptivity of the product of the first slow step of the reaction is 176. Also, at a pH higher than 6, a blue complex having a smaller molar absorptivity than the purple con~plexbegins to predominate in the solution.
' HAMM, R. C., J. Am. Chem. Soc., 75,5670 (1953).
For the purposes of this experiment, the reaction conditions are limited to a pH range between 3.5 and 5.5, where the purple complex predominates. The reaction is run at room temperature, and the student is asked to find the reaction order exponents for the hypothetical reaction expression: Rate = -d[Cr(III)I - +dlCr(lII)-EDTAI = k,Cr(III)I.IH+lb (,) dt
dt
To simplify matters, the concentration of EDTA is kept high (0.1 M) with respect to chromium(II1) so that its concentration can be neglected in the mathematical rate expression. Actually, Hamm found that the rate of the reaction did not depend upon the concentration of EDTA. Several containers of 0.1 IM EDTA are adjusted to different pH values using coucentrated sodium hydroxide or nitric arid. At t,ime = 0, a small aliquot of concentrated chromium(II1) nitrate solution is added to each container, so that the final chrominm(II1) concentration is approximately 6 X molar. I t is not necessary to measure concentrations or volumes exactly, or to run a "blank," because absolute rate constants for the reaction will not be measured. Data is taken over a four-hour laboratory period. At the end of t,he experiment, the solutions are heated to 100°C in a water bath for 10 minutes to complete the reactions, and the absorbance of each solution at time = infinity is measured. This. data is then used to calculate an absorbance value which is a measure of the unreacted chromium(II1) :
a RINGBOM, A,, "Complexatition in Analytical Chemistly,"Interscience Publishen, he., New York, 1963.
When the quantity A , - A, is plotted against time.
Figure 1. Absorption spectrum of the chrornivmlllll-EDTA complex a t pH 4.1 meorwed using a Spectronic-20 spectrophotometer.
Figure 2. Plot of A, - Ar versus time in minuter for ~everolpH valuer. The large circles represent doto used to construct curves which show the. rate os a function of the concentration of each reoctmt.
Volume 42, Number 9, September 1965
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479
-1.01 30
-OD
40 -05 upper scale:-log,,~H*l
50
-10
60 -I. 5
I time, rniwter, upper ICOIO
Lower scole: + logo(A--At1
Figure 3. When loglratel is plotted against log(Crlll1 or -iog(Htl, straight liner are obtoined which hove slopes q u a i to tho experimentd re~ctionorder exponents for the re3pective reactant,.
for different pH values, a family of curves is obtained which is used to obtain the reaction order exponents. An example of this dat,a plot is shown in Figure 2. To show the student exactly how rate data is processed, points are chosen on the curves under reaction conditions representing constant pH values and constant values of the chromium(I1I) ion concentration. This is shown in Figure 3 as circled data points. The slope at each point is measured by a graphical method, and the rate of the 'eaction at that point is estimated. Plots are made of the logarithm of the rate versus the logarithm of the chosen reactant concentration while the other reactant is held constant. An example of this type of curve is shown in Figure 3. The slope of each log-log plot gives the reaction order exponent. The table shows the values obtained by several students. If the reaction order exponents are unity, the reaction data can be tested for conformity to pseudo firstorder behavior by plotting the logarithm of the concentration of chromium(I11) or hydrogen ions versus time and noting the linearity of the plot. An example of this treatment is shown in Figure 4. The reaction beheen chromium(II1) and EDTA
pH 5 8
Figure 4. The reaction order exponent for chromiurnlilll is close to unity. This plot shows thot indeed pseudo fist-order kinetics is followed by tho reoctont when the pH is held conrtont.
provides a useful vehicle for several interesting laboratory experiments in chemical kinetics for different fields of chemistry. Determination of the activation energy of the complex, establishing exact mechanisms using extended reaction conditions, and investigating the catalytic effect of different substances are suggestions for extending the reaction to other experiments. Exuerimentol Reoction Order Exuonents
Source
Order in Cr(II1)
Order in iH+i
0.90
-0.90
1.10
-1.11
VH
Student
Author Student I1 111 IV
5.8
v
VI Average Average value.
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480
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Journal of Chemical Education
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