A Kinetic Experiment Using Potentiometric
R. H. Smith Mocquorie university North Ryde, N.S.W., 2113, Australia
Determination of Reactant concentration
Undergraduate kinetic experiments would appear to have three major aims: (1) to provide experience in obtaining and processing kinetic data, (2) to give an idea of the scope and variety of experimental methods which can be used for rate measurements, and (3) to illustrate how kinetic results can provide information about the mechanism of a reaction. Frequently however published experiments do not achieve all these aims, (3) generally being the one which is neglected. This report describes an experiment which appears to satisfy all three aims and which in the author's experience works extremely well, either as a formal class experiment or as a project type exercise of longer duration. Essentially i t involves the measurement of the rate of a reaction by a potentiometric technique, and utilizes the fact that, if the sensing electrode responds to the concentration of a reactant and if the reaction is first order (or pseudo first order) in that reactant, then the emf will be linear with time. The electrode used is a bromine electrode which is used in a manner based upon that described by Bell and Ramsden.l The reactions used are oxidation of formate and oxalate by bromine HCOO-
+
Br,
+
C02
+ H' +
2Br-
I t is very easy to convince students that there are several possible mechanisms for each of these reactions and that all of them look chemically reasonable: for example, for the formate reaction some possible rate determining steps (preceded by the necessary rapid pre-equilibria) are
Br-
+
HBr
+
CO*
+
mechanistically. A pair of kinetically indistinguishable mechanisms can also be included if desired. In order to simplify kinetic analysis the initial concentrations of HCOOH, H+, and Br- are made much greater than that of bromine and hence remain essentially constant throughout the reaction. Thus a generalized rate law Rate = ~ [ H C O O H I ' [ H ' ~ [ B ~ - ~ [ B ~ ~
simplifies to Rate = kn[BrJ" where k,, the pseudo nth order rate constant, is given by
By determining which function of [BIZ] is linear with time ([Brz], log @rz], l/[Brz] etc.) one can determine n and hence a value for k, from that experiment. Alternatively if all the proposed mechanisms lead to rate laws which are first-order in bromine we can take as our working hypothesis that Rate = kdBr,l and simply test whether or not this is true. Then by repeating the experimeht with a different (but still excessive) initial concentration of formic acid and seeing how k , changes as [HCOOH] is altered, a value of a is obtained. Similarly b and c are determined. The experiment has been used in two different ways: first as a normal class experiment with students working in pairs, in which form it has usually been possible to per-
(1)
'02 By having to derive the rate law implied by each of several proposed mechanisms, students are easily persuaded that a determination of the rate law will lead to elimination of many of the possibilities, and thus before starting the experiment they can see how i t will he of value P., and Ramsden, E.N.,J. Chem. Soc., 161 (1958). Smith, R. H.,A,ut.J. Chem., 25,2503 (1972).
1 Bell, R.
Figure 1. The thermostatted reaction vessel. A, Saturated calomel electrode: 6, platinum electrode; C, magnetic stirrer; D, thermostatting water inlet.
Volume 50, Number 6, June 7973 / 441
form the minimum number of experiments (usually 4-6) to distinguish between the four mechanisms just described in about 6-9 hr (including time for processing results), providing stock solutions of t h e reagents are available. Formate was t h e only reductant used in this form of the experiment. Secondly the experiment has been used i n a more extended form a s a "project" of about 25 hr duration for pairs of students: both oxalate and formate have been used a s the reductant i n this form. Such pairs of students (formate t o one pair, oxalate t o t h e other) have been able to establish t h e rate laws more thoroughly in this time and for t h e formate reaction (which is considerably simpler t o study) the activation energy bas also been measured (using three temperatures). T h e fact t h a t these a p parently similar reactions have different rate laws2 has been the cause of keen competition between pairs of students (who assume t h e reactions have analogous rate laws) a s they try t o prove the correctness of their particular rate laws. Since t h e time to perform the actual kinetic experiment is not great (10-20 min), one set of apparatus per pair of students is adequate. Experimental
The apparatus, shown in Figure 1, consists of a thermostatted reaction vessel sitting upon a magnetic stirrer. The vessel is made with a B55 standard socket at the top into which fits the lid (made from a B55 cone) whieh carries a platinum wire electrode and a commercial saturated calomel electrode (electrical cannection via a sintered glass plug). The vessel used for the oxalate reaction is painted black to eaclude light completely: for formate a transparent vessel is used since the normal levels of laboratory light do not affect the rate. The formate reaction is thus much better for a normal student experiment since the student can see what is occurring. The platinum electrode needs occasional cleaning. This is most simolv done bv heatine the wire to red heat for about a minute at the'bknnine-of eachhav of the exoeriment. The reaction mixture nnmt bestirred at a un~f
