Martyn V . Twigg Erindale College Universitv of Toronto Toronto 181, Canada
II I
Aquation of Trb-(l,lO-phenanthr~he)iron(II) in Acid Solution A kinetics experiment
During recent years, kinetic studies have produced much information about the mechanisms by which chemical reactions occur. The continued interest in reaction kinetics is reflected in the content of undergraduate courses-this topic is now part of the first year curriculum at most universities, and frequently some study of the subject is made at high school. However, a parallel development of the topic in the laboratory is hampered by the extended period of time or involved procedures needed to follow the progress of most of the usual, rather slow, reactions that are used to illustrate various rate laws. As a result it is tedious, as well as time-consuming, to determine rate constants a t several temperatures to enable the activation energy of the reaction to be determined. Nonetheless a t the introductory level the practical aspects of reaction kinetics, and in particular familiarization with the exponential temperature coefficient of chemical reactions, are important. Therefore a search of the literature was made to find a suitable reaction whose activation energy could be determined by a student within a normal 3-hr laboratory period. Such an experiment has to he comparatively simple, yet illustrate the fundamental concepts of reaction kinetics. The aquation of the tris-(1,lO-phenanthro1ine)iron(11) cation in dilute mineral acid was considered suitable, since i t is a reaction for which reliable kinetic data are available (1-5) and it has an easily measured rate a t accessible temperatures: also, only small quantities of reactant are required if advantage is taken of its intense color to follow the course of the reaction. Further, the reactant complex is easily prepared in solution, and the only other reagent required is a dilute solution of mineral acid. Moreover, the intense color of the complex suggested that a simple visual comparison technique could he used to determine the rate constants at different temperatures, as well as the obvious spectrophotometric method which could he employed if suitable instruments are available. WphenX"
-!%!?+ phen =
Fe(H*0Ie2+
+
&
3phenHC
Tris-(1,lO-phenauthroline)iron(II) is a very stable (& = lo2' (4)) 8- low-spin substitution inert complex. However, in all but extremely dilute solutions of mineral acid, protonation of the free ligand causes the complex to dissociate in an attempt to restore equilibrium conditions with the free ligand. Finally, none
of the complex remains and the resulting solution is colorless. The reaction is first order in complex, loss of the first ligand is rate determining (2) and a t a given temperature the rate is not markedly dependent on the acid concentration (2, 3 ) . Addition of salts has been shown (8) to produce a small and sometimes significant effect on the observed rate constant. This paper describes bow first-year students, working in pairs, successfully studied this reaction using an elementary spectrophotometer and an extremely simple visual half-life method to determine rate constants a t various temperatures, so enabling the activation energy of the reaction to be determined. The Experiment The necessary stock solution of IFe(phen)slSO, is prepared by dissolving pure ferrous ammonium sulfate (1.32 g) and 1,lOphenanthroline hydrate (2.0 g) in 300 ml of distilled water. It is necessary to warm (50°C) and stir the mixture for complete dissolution of the ligand. This quantity of solution is sufficient for about one hundred students working in pairs-once prepared it has good keeping properties. A stock solution of dilute mineral acid is also required1-2 M hydrochloric acid ("bench strength") has been used with success and solutions down to 0.5 M can be used. Procedure A:
The Speciroscopic Method
Sufficient of the ferrous complex solution (about 6 drops) is added to 45 ml of the stock acid solution contained in a 50-ml conical flask, so that at the wavelength of maximum absorption in the visible region (6) 510 nm, the resulting solution has an optical density of about one in the cell being used. The stoppered flask is suspended in a thermostat bath, the time noted and after about 4 min a sample (-4 ml, but volume depends on the type of cell being used) is pipetted from the flask. The time at which the sample is taken is noted and its optical density at 510 nm quickly determined before it is discarded. This procedure is repeated periodically until tbe optical density is less than 0.2 or more t h n 2 hr have elapsed. The time between taking samples is determined by the temperature of the bath: 25°C 15-20 min, 30°C -10 min, 3 5 T -5 min, 40°C -3 min. A table of time (min), optical density (OD), and log OD is made and the rate constant at the temperature concerned obtained from the slope of 8 graph of log OD versus time.= Procedure B:
The Visuol Holf-life Method
In this procedure, the half-Me of the reaction is measured directly by a. visual comparison technique. A beaker (400 ml is ideal) containing a, magnetic follower is three quarters filled
' Nitric and perchloric acids are not recommended. Nitric acid oxidizes the iron to the ferric state and perchloric acid precipitates the complex cation as the sparingly soluble perchlorate sdt. Because the optical density of the reacting solution ultimately becomes zero the first-order rate plot is simply In OD versus time. Volume 49, Number 5, May 1972
/
371
with water and slowly heated to about 42'C on a magnetic stirrer heater. Throughout each "kinetic mn" the temperature of the water, measured with a. 0-llO°C thermometer, is kept constant to better than 10.5'C by careful control of the rate of heating and, when necessary, the addition of smdl quantities of cold water from a wash bottle. Two similar test tubes are immersed in the water. One, the standard, contains 9.50 ml of distilled water and 0.50 ml of the stock f e m u s complex solution (measured with a 1 ml disposable syringe). The other test tube, initially containing 9.00 ml of stock acid solution, (conveniently measured with a buret), is also placed in the beaker. After a few minutes when the temperature of the solutions in the test tubes is that of the surrounding water, 1.00 ml of the f e m u s complex solution is quickly injected from a 1-ml disposable syringe into the acid solution. When injection is rapid, mixing t a k a place quickly without the test tube having to be removed from the beaker. At the time of injection a stop watch is started, whiehis stopped when the colors of the two solutions are judged to he the same. This measured time is the half-life ( t t / , ) of the reaction a t the temperature concerned. The temperature is then increased by about 4"C, and using the same standard, hut a fresh portion of acid, the half-life is determined at the new temperature. This procedure is repeated until the half-life is about one minute (about 60DC). The rate constant a t each temperature is calculated from the relation k = 0.693/11/, and the activation energy obtain& from an Arrhenius plot.
Results
Several pairs of first year students, having better than average capability at laboratory work, measured the rate of aquation of tris-(1,lO-phenanthroline)iron(II) in 2 M hydrochloric acid at four temperatures in the range 25-40°C by procedure A, using a Spectronic 20 spectrophotometer, and at higher temperatures with the half-life method, to check the agreement of the data obtained from the two methods, as judged by the linearity of the resultant Arrhenius plot.3 Figure 1
Figure 2. 3 life method.
f log kos. versus l/ToA.
Squwes, rote plot; circler, half-
and a t six and seven different temperatures in the range 40-60°C by the visual half-life method (Procedure B). This was easily completed within a single 3-hr laboratory period, and most students obtained good results. Since it is known that the rate of the reaction decreases with increasing acid concentration (2), the average value obtained for therate constant at 30°C in 2 M HCl, 9.2 X min-', is in reasonable agreement with the published (1) value of 9.8 X lo-%min-' obtained in 1.07 N HBO,. Generally, data obtained from procedure B produced fairly good Arrhenius plots, and normally little trouble was encountered in judging when the two color intensities were the same. Occasionally the data point from the rate plot was obviously not within experimental error of the best straight line through the points on the Arrhenius plot. This could be due to consistently misjudging the time at which the intensities of the two solutions are the same, or more likely to the reference solution being incorrectly prepared. Errors in making up this solution cause the best straight line on the Arrhenius plot to be displaced up or down without changing the s l o p e . ' T h e calculated activation energy was usually within 2 kcal of the literature value. The students who performed these trial experiments reported that they enjoyed them because the objectives were clear and the practical work, as well as the calculations, was not difficult, yet the experiment illustrated important aspects of reaction kinetics covered in lectures. Conclusions
+
MINUTES
log OD verrur time for the aquation of Fe(phenhz+ Figure 1. Plots of 1 in hydrochloric m i d ot various temperatures (dato of E. J. Greengloss ond
L J . 1. Honlon).
shows typical rate plots obtained a t different temperatures. Rate constants obtained from these graphs are combined with those from the half-life method to give the Arrhenius plot shown in Figure 2. The agreement between the two sets of data is good, and the calculated activation energy (29.6 kcal) is in very good agreement with the literature value, (30.0 0.2 kcal) (1) showing that both methods can produce reliable results. A larger group of first year students, about sixty pairs, then, as part of their normal laboratory work, measured the rate of the aquation of tris-(1,lO-phenanthroline)iron(II) in 2 M hydrochloric acid at 30°C using a "Spectronic 20" spectrophotometer in procedure A
*
372
/
Journal o f Chemical Education
The results obtained from this series of trial experiments indicate that meaningful results are easily obtained using a simple half-life method, enabling the students to "see" both qualitatively and quantitatively the effect of temperature on the aquation of tris-(1,lOphenanthroline)iron(II) in acid solution. If spectrophotometers and thermostat baths are available, good rate plots for this reaction are easily obtained, and careful workers quickly obtain an excellent Arrhenius plot. The acid concentrations in procedures A and B are not exactly the same. This produces only a very smdl systematic difference in the rate constants obtained by the two methods. ' The volume of the ferrous complex solution used in the reference solution should he carefully measured. E m r s in its measurement produce a systematic error in the measured rate constant equal to approximately one and a half times the oliginal percentage error, but in the opposite sense.
Acknowledgment
Literature Cited
Without the interest and enthusiasm of the parti& pating students of the 197G71 Erindale College CHM 120 E class, this investigation no+, have been possible.
(ll Bmaess. I.,*ND PRINCD,R. H., J . Cham. Sm., 5752 (1963). J. E.,R*~oLo. F..ANDNEUMIINN, H. M., J . Amcr. Chem. Soc., (2) DICKBNS, 79, 1286 (1957). (3) L=E.T.% KKOLTHOPF. I. M., A N D L E ~ W ND. ~ ,L.,J. A m , . Cham. Sac.. 70, 3596 (1948). (4) S ~ I I E N G. . L.,A N D MARTELL. A. E., " ~ t a h i i i t yconstants of ~etal- on Complexes" (2nd ed.). The Chemical Society. London. 1964, p:665. (5) S c n r ~A. ~ ,A,. "The Analytled Applioatlona of ],lo-Phenanthroime and Related Compounds," Pergamon Preas, New York. 1969.
Volume 49, Number 5, May 1972
/
373
