Peter Jones. and K. B. Oldham King's College, University of Durham' Newcastle upon Tyne, England
I
I
The theory o f the
Formaldehyde Clork b a d i o n
The "formaldehyde clock" reaction has been recommended, both in THIS JOURNAL^ and the School Science Review3 as an experiment suitable for lecture demonstration and for students to perform. It is indeed a "tested" demonstration.' A theory of the reaction, first given by Wagners has been accepted by later authors. Our purpose is, first, to point out that Wagner's theory is quite unsound and, secondly, to give a more satisfactory explauation. According to Wagner, the following reactions occur: HCHO
H20
OH-
---
+ HSOa-
+ HCHO + SOa?+ HSOa-
CHz(OH)SOs-
CH2(0H)SOsSOsz-
+ HIO
+ OH-
(1) (2)
(3)
Reactions (1) and (2) are regarded as rate-determiniu~g, reaction (3) as instantaneous. Reaction (2) is introduced to account for the observation that the reaction time r (the time taken from initiating the reaction to observing an indicator color change) is decreased by increasing the initial sulfite concentration a t constant initial formaldehyde and bisulfite concentrations. I t is considered that the pH increases only when all the original bisulfite is used up. Our objections are as follows. (1) The theory supposes that a chemical reaction can go to completion in a Jnite time which can be altered by changing initial concentrations; this is clearly wrong. (2) The theory ignores the buffer action of the sulfite/bisulfite system:
where Kinis the dissociation constant of the indicator and (CR) is the indicator color ratio at which the end point is detected by an observer using a particular indicator concentration. For simplicity of treatment we will assume that [H+], is determined by the initial stoichiometric buffer ratio and that [SOa2-] remains constant up to time r, i.e., [Hf]- = KJHS0,-],/[SO,B-1,
A rigorous solution leads to the same result as that obtained below, for the conditions used by Wagner. For initial concentrations of formaldehyde (FO), bisnlfite (R,)and sulfite (SO)such that Fo > BO the integrated second order rate equation between the limitst = Oandt = is:
Making the appropriate substitution yields:
-
Since for phenolpl~thaleinKi.'/K, 10-5 SO(Kinl/K.)in Wagner's experiment.
((lo -
R,,) >>
Introducing this simplification, (8) rearranges to
In the graph Wagner's original data are compared with equation (9). This yields the values at 25OC. We show below that, if Ii, = 6.7 X this equilibrium is taken into account, the effect of changing sulfite concentration is accounted for without involving the improbable termolecular process equation (2). We suggest that the kinetics of the formaldehyde clock reaction can be explained simply in terms of reaction (1) and the rapid equilibrium (4). Reaction (1) is considered to be irreversible. The reaction time ( T ) then corresponds to the time taken for reaction (1) to produce a change in [Hf] from an initial value,
to a value a t time r given by [Ht],
=
Ki.(CR) = K i n '
(6)
On August 1, 1963, King's College will become the Universit,y Newcastle upon Tyne. R. L., J. CHEM.EDUC.,32.78-9 (195.5). RARRETT, "HISMAN, D. G.,Seh. Sn'. Rev., 38, lO&l (1956). ' ALYEA,H. N., A N D I~IITTON, F. B., eds., "Tested ilemrmst,rntions in Chemistry," 4th ed., Chemical Education Publishinr: Co., Easton, Ps., 1960, p. 59. j WAGNER, C., Rer., 62, 2873-7 (1927).
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/
Journal of Chemical Education
k
=
8.9 liter rnule-'sec-~ and (pK,,'
- pK,)
=
1.9
There is some spread of the results and this is hardly surprising when the following sources of error are coo-
sidered. (1)There are timing errors if T is small. (2) There are timimg errors if, a t the endpoint, dpH/dt is small. We have followed the reaction using indicator and a recording pH meter simultaneously and have shown that the phenolphthalein end point does not always occur in a region of rapidly changing pH. (3) In Wagner's work the ionic strength varied considerably between experiments (from about 0.1 to 1) so that K., K,,' and probably kl would not remain constant. (4) Autoxidation of sulfite to sulfate may cause an increase in the apparent r , or the end point may not appear, or it may appear and rapidly disappear (competition from an "oxygen clock"). The three points enclosed by a rectangle in the graph indicate the reproducibility Wagner obtained under a single set of conditions.
Barrettz found a linear relation between .r and dilu[So]mere kept constant. tion when the ratios [Fa]:[Bo]: We do not understand his explanation hut this result is readily accounted for by equation (9). Barrettz also gives an empirical linear relationship between temperature (in OC) and log r. We do not regard this as a useful student exercise since the empirical relation has no theoretical significance, and the function actually describing the temperature coefficient of T is complex, involving the temperature coefficients of kl, K. and Kt,'. Acknowledgment. We wish to thank Dr. A. K. Covington for the use of the recording pH meter and Dr. M. L. Haggett for experimental assistance.
Volume 40, Number 7, July 1963
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