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Page 1: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

616 PERIODIC CHEMICAL REACTION

A PERIODIC CHEMICAL REACTION

THE REACTION BETWEEN HYDROGEN PEROXIDE AND IODIC ACID

BY M. G. PEARD AND C . F . CULLIS

Received 6th November, 1950

In the reaction between iodic acid and hydrogen peroxide, the rate of evolu- tion of oxygen may, in certain circumstances, vary periodically with time. Factors affecting the periodicity have been examined, and the mechanism of the reaction is discussed. It is suggested that the periodicity is the result of the interplay of two factors, the autocatalytic nature of the reaction, and the physical removal of the product responsible for the catalysis by the other product, oxygen.

During the last hundred years, many miscellaneous examples of periodicity in chemical systems have been discovered,' among which not the least interesting are those in which the velocity of a chemical reaction varies periodically with time.

The majority of such reactions at present known appear to fall into one of two classes :

(I) Reactions in which gases are liberated as a result of chemical reaction

This class is exemplified by the work of Ostwald and of Hedges and on the dissolution of metals in acids, where periodic variations in

The ex-

(2) Homogeneous vapour-phase oxidations involving a number of auto-

One of the earliest examples is the oxidation of hydrogen ~ulphide,~ in which under certain conditions reaction is accompanied by a series of well-defined explosions. Newitt and Thornes observed a similar periodic occurrence of " cool flames " during the slow combustion of some higher paraffins, and analogous beliaviour has been found in the oxidation of

For reviews, see Hedges and Myers, The Problem of Physicochemical Perio- dicity (Arnold, 1926) ; Veil, Actualitbs Scientifiques et Industriels (I934), No. 154.

Ostwald, Physik. 2.. 1899, 21, 88. Hedges and Myers, J. Chem. SOC., 1924, 604. Thompson, Nature, 1931, 127, 629 ; Thompson and Iielland, J. Chem. Soc.,

Newitt and Thornes, ibid., 1937, 165G.

at a solid-liquid interface.

Myers the rate of hydrogen evolution may frequently be observed. planation of this type of periodicity is probably largely physical.

catalytic stages.

1931, 1809.

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Page 2: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARD AND C. F. CULLIS 617 other aliphatic compounds.6 Explanations of the periodic nature of such reactions have been suggested by Frank-Kamenetzky

There is one periodic reaction, however, that does not belong to either of these categories, namely, the reaction between iodic acid and hydrogen peroxide. This is an apparently homogeneous reaction in solution in which the rate in certain circumstances varies periodically with time. The periodicity was discovered by Bray,* who supposed that the reaction proceeds in two main stages each of which is undoubtedly complex :

and Salnikov.8

5 Hz02 + 2 HIO, = 5 0, + 1, + 6 H,O 5 H,O, + I, = 2 HIO, + 4 H,O

10 H,O, = 10 H,O + 5 0, The net result is therefore the catalytic decomposition of hydrogen per- oxide. Bray found that between fairly narrow limits of hydrogen ion concentration the rate of oxygen evolution varies periodically as the reaction proceeds, but he did not advance any precise explanation of the curious results obtained. It was considered of interest by the present authors to carry out a detailed study of the reaction, and to attempt to elucidate the nature and cause of the periodicity.

Experimental The reaction between potassium iodate, hydrogen peroxide and sulphuric

acid a t 6ooC proceeds with the liberation of oxygen and small quantities of iodine. Various criteria may be used to follow the course of the reaction.

(1) Measurement of Rate of Oxygen Evolution.-The majority of results were obtained by measurement with a Lunge nitrometer of the volume of oxygen evolved from the reaction mixture. Even at 60' C oxygen tends to form supersaturated solutions. The work of Richardson and Bailey lo shows how inefficient rotatory stirring can be as a means of eliminating supersaturation, so that i t was considered desirable to use another method.

FIG. I.

13 Bardwell and Hinshelwood, Proc. Roy. SOC. A , 1951,205, 375. 7 Frank-Kamenetzky, J. Physic. Cham. (U.S.S.R.), 1940. 14, I ; Walsh,

8 Salnikov, Compt. rend. U.S.S.R., 1948, 60, 405, 611. Q Bray, J. Amer. Chem. SOC., 1921, 43, 1262. 10 Richardson and Bailey, Proc. Roy. SOC. (Dublin), 1934, 21, 43.

Trans. Faraday SOC., 1947, 43, 305.

23

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Page 3: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

61 8 PERIODIC CHEMICAL REACTION If a flask be supported by a

slightly flexible mount and struck sharply at regular intervals, the resulting passage of shock waves and general agitation brings about effective de-super- saturation of the liquid therein.

The apparatus constructed for this purpose (Fig. I) consists of an electric motor directly connected to a reduction gear, the axle of which carries a brass wheel W of diameter 3 in. A short pin with ball-race B attached was fastened through the edge of the wheel. The ball-race makes contact with a short piece of metal M screwed to the side of the rod R and, as the wheel rotates, pushes the rod against the tension of the spring S until the end of the projection is reached. At this point the spring pulls the rod back sharply, and the rubber bung at- tached to the lower end of the latter strikes the flask, which is supported by a strong steel spring. It was found that satisfactory results were obtained if, with considerable spring tension, the flask was struck approximately once a second.

The method adopted was that due to Bell.”

FIG. 2 .-The on the course of reaction a t 60° C.

[KIOJ : 0.100 M ; [H20Z] : 0.226 M : [H2S0,] : 0.0767 N. a: Agitation of reaction mixture by vigorous shaking with Vibro agitator

(A. R. Parsons & Co.). b : Agitation of reaction mixture by method described in Fig. I . c : No agitation of reaction mixture.

The influence of efficient agitation of the reaction mixture may be seen from Fig. 2, which shows that the removal of supersaturation reveals the periodic nature of the reaction. Curve a, obtained by using another means of agitation, is provided for comparison. Readings of the volume of oxygen evolved were taken at intervals of I rnin.

(2) Measurement of Iodine Concentration.-The iodine concentration during the reaction is found to be dependent on the rate of oxygen evolution. This is due to the fact that iodine is appreciably volatile at 60’ C, so that the oxygen evolved from the reaction mixture tends to carry with it some of the iodine (see later). Unless supersaturation is eliminated by suitable agitation of the reaction mixture, measurements of the iodine concentration are of little significance because the correct conditions for periodic evolution of oxygen are not achieved. Contin- uous measurement of the iodine concentration by physical means (e.g. colorimetric) is thus made difficult. Samples were therefore withdrawn from the reaction mix- ture a t frequent intervals and the iodine estimated by extraction with carbon tetrachloride and comparison with a standard in a Duboscq colorimeter.1

l1 Bell and Baughan, Proc. Roy. SOC. A , 1937, 158, 464 : Bell and Trotman- Dickenson, J. Chem. SOC., 1949, 1288.

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Page 4: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

11.1. G. PEARD AND C. F. CULLIS 619 (3) Measurement of the Hydrogen Ion Concentration.-The hydrogen

ion concentration was measured continuously with a glass electrode and pH meter. The glass electrode was connected to the calomel electrode of the pH meter by a bridge of strong ammonium nitrate solution. The reaction mixture inevitably becomes contaminated with a little ammonium nitrate. Control experiments were ca,rried out with various quantities of ammonium nitrate. The rate of reaction is slightly depressed but the effect is not of quantitative significance. Since the calomel electrode was a t room temperature and the glass olectrode was at 60° C, relative rather than absolute pH values were obtained.

Results The Course of the Reaction.-Studies of the rate of oxygen evolution

provided most of the experimental data. The results obtained are fairly repro- ducible, though they become less so at high frequencies of periodicity. The rate of reaction is sensitive to traces of chloride ion and to the concentration of iodine present. The latter tends to be dected by slight variations in the rate of agita- tion of the reaction mixture since, as already mentioned, iodine is appreciably volatile a t 60° C.

The general course of the reaction with respect to oxygen evolution is shown in Fig. 2. The variations of iodine and hydrogen ion concentrations during a typical experiment may be seen in Fig. 3. There is a periodic variation in the concentration of iodine, visual observation of the reaction mixture showing that its value reaches a maximum just before the occurrence of a maximum in the rate of oxygen evolution ; whereupon the brown colour of the solution fades, and immediately afterwards the space above the liquid is seen to be deeply coloured by iodine vapour. The hydrogen ion concentration, on the other hand, remains effectively constant (to within 0.02 pH unit) throughout the reaction.

I . / # M/j7ute.s 20 40 60

FIG. 3.-The variation of iodine and hydrogen ion concentrations during re- action at 60' C.

[KIO,] : 0.100 M : [HBOJ : 0.228 M ; [H,S04] : 0.0958 N.

/

a : [I2] (moles/l. x 104) ; b : pH.

The Stoichiometry of the Reaction.-According to the equations given earlier, the net result of the reaction is the decomposition of hydrogen peroxide, the potassium iodate and sulphuric scid apparently remaining in their original concentrations after reaction is complete. Since iodine is readily lost by vapor- ization from the solution there is, however, found to be an appreciable decrease in iodate concentration at the end of the reaction. This decrease was estimated by allowing the iodate present to react with hydrogen iodide and subsequently titrating the liberated iodine with thiosdphate, a typical result being shown in Table I. From this i t may be seen that 5-9 yo of the iodate initially present has been decomposed at the end of the reaction.

TABLE I (Na,S,O, = 0.548 N)

Initial titre = 54-4 ml. = 0.0149 mole iodine. Final titre = 51-2 ml. ZE 0.0140 mole iodine.

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Page 5: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

620 PERIODIC CHEMICAL REACTION The fact that the reaction does not involve decomposition of hydrogen per-

oxide only, may also be deduced from the volume of oxygen liberated. 50 ml. 0.226 M H,O, should liberate 126.6 ml. oxygen at N.T.P., whereas 132.0 ml. are in fact evolved from this quantity of hydrogen peroxide. The balance (5.4 ml.) must presumably have come from iodate that has been irreversibly reduced to iodine :

5 H,O, + 2 HIO, -+ 5 0, + I, + 6 H20. It can be calculated on this basis that 3-9 yo of the iodate initially present has been decomposed. The iodate has probably been lost by vaporization out of solution as iodine.

FIG. +-The influence of hydro- gen ion concentration on the

course of reaction at 60' C. [KIO,] : 0.100 M ; [H20,] : 0.226 M.

CH~S0,l : a = 0.123 N ; b = 0-113 N ; c = 0.0958 N ; d = 0.0767 N ; e = 0.0740 N ; f = 0.0682 N.

FIG. 5.-The inf lu- ence of reactant con- centrations on the average rate of re-

action at 60' C. 0 [H,O,]: 0-226M:

[H,SO,]: oe117N; [KIO,] varied. [KIO,] : 0.100 M; [HeOz] : 0.226 M ; [H,SO,] vaned. [KIO,] : 0-IOOM; [H,SO,] : 0.0958 N; [H,O,] varied.

Influence of Hydrogen Ion Concentration.-Fig. 4 shows the effect of varying concentrations of sulphuric acid on the rate of oxygen evolution. Similar results were obtained by Bray.s The apparent order with respect to hydrogen ion may be deduced from Fig. 5 in which log,, [H+] is pIotted against log,, (average rate of reaction) ; its value is found to be approximately 2. In Fig. 6

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Page 6: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARD AND C. F. CULL'IS 62 I

is plotted the number of periods observed during a fixed time against the square of the acid concentration, a linear relationship being obtained. Though the significance of this is not clear, the result indicates the sensitivity of the perio- dicity to the hydrogen ion concentration.

FIG. 6.-The influence of reactant concentrations on the frequency of

periodicity at 60' C. 0 [H,02] : 0 2 2 6 M ;

[H,SO,] : 0.117 N ; [KIO,] varied. [KIO,] : O*IOO M ; [H,02] : 0.226 M ; [H,SO,] varied.

4' 8' u 12' /

Influence of Iodate Concentration.-The effect of varying concentrations of potassium iodate on the rate of reaction is shown in Fig. 7. The apparent order with respect to iodate (Fig. 5 ) is approximately unity and a linear relation- ship also exists between the frequency of periodicity and the iodate concentration (Fig. 6).

FIG. 7.-The influence of iodate concentration on the course of reaction a t 60' C. [H202] : 0.226 M ; [H,SO,] : 0.117 N.

[KIOJ : a = 1-136 M ; b = 0.100 M ; c = 0.080 M ; d = 0.060 M ; e = 0.040M-

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Page 7: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

622 PERIODIC CHEMICAL REACTION Influence of Hydrogen Peroxide Concentration.-The effect on the

rate of oxygen evolution of altering the hydrogen peroxide concentration is shown in Fig. 8, from which i t is seen that the frequency of periodicity is little altered by such variations. Fig. 5 shows that the apparent order with respect to hydrogen peroxide is approximately unity.

FIG. &-The influence of hydrogen peroxide concentration on the course d reaction at 60' C.

[KIOJ : 0-100 PA ; [H,SO,] : 0.0958 N. [H,O,]i: a = 0.339 M ; b = 0.282 M ; c = 0-226 M ; d = 0.169 M ; e = 0.113 M.

It appears therefore that the rate of oxygen evolution is related to the reactant concentrations by an expression of the form

(d[O,I/dt)a"era,e~ = ICH +I"IO,-l CHnOe3 . Influence of Temperature.-The influence of temperature on the reaction

is shown in Fig. 9. The rate of oxygen evolution increases fairly rapidly with temperature, but the periodic nature of the reaction persists over a range of more than zoo C. From the results given i t is found that the temperature coefficient corresponds to an apparent activation energy of 23 f 2 kcal./mole.

Influence of Surface.-A large increase in the effective surface area of the reaction vessel (obtained by filling the vessel with glass beads) causes a small decrease in the rate of reaction and the frequency of periodicity. The effect may perhaps be due to the greater surface of reaction mixture exposed facilitating loss of iodine from the solution by vaporization.

Influence of Iodine.-Fig. 10 shows the effect of adding various quantities of iodine to the reaction mixture. Curve a shows that iodine in appreciable quantities strongly catalyzes the decomposition of hydrogen peroxide in presence of iodate. Smaller quantities accelerate the rate of reaction in the initial stages. Parallel experiments in the absence of iodate (Fig. 11) show that iodine alone in acid solution does not appreciably catalyze the decomposition of hydrogen peroxide. Curve e was obtained by measuring the rate of oxygen evolution when a layer of carbon tetrachloride was present to extract the iodine as fast as i t was formed.

The iodine concentration thus seems to be a factor of great importance in the reaction. Furthermore, since iodine is rather volatile at 6 o o C , the iodine concentration will be affected by the surface area of the liquid-gas interface and hence the rate of agitation. For this reason, the frequency of periodicity is not always perfectly reproducible.

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Page 8: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARD AND C. F. CULLIS 623

rn v t

- 1 40 5

91

- 3

1 120 140 f 160 FIG. 9.-The influence of temperature on the

course of reaction.

psSO,] : 0.0862 N.

a = 77.3" C ; b = 7 2 0 3 ~ C ; c = 66.7' C ; d = 63.5' C; e = 60.0~ C;] f = 57.2' C.

[KIO,] : O'IOO M; m,OJ : 0.226 M ;

Temperature :

FIG. 10.-The influence of added iodine on the course of the reaction at 60° C.

(KIOJ : 0.100 M ; [H,OJ : 0.228 M ; @,SO,] 0.0876 N.

[Id : u = saturated (< 0.0245 M) : b = 0.0046M; c = 0.0019 M ; d = no added iodine: e = no added iodine; reaction mix-

ture in contact with a layer of carbon tetrachloride.

of hydrogen peroxide at 60' C. a : [HeOe] = 0.226 M ; [Cl-] = 2.0 M. b : [H20e] = 0'226M; [H,SO,] = 0-086N [Iz] = 0-0015 M.

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Page 9: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

624 PERIODIC CHEMICAL REACTION

powerful influence on the reaction rate. Curve f indicates that such ions catalyze the reaction

2 HIO3 + 5 H,O, 3 5 0, + I, + 6 H,O. The frequency of periodicity (as measured by the difference in the number of periods with and without varying additions of chloride ion) is found to be ap- proximately proportional to the square of the chloride ion concentration. Some comparative results in absence of iodate are given in Fig. 11 and i t will be seen from these that the effects shown in Fig. 12 cannot be attributed to the catalytic effect of chloride ions in decomposing hydrogen peroxide alone.

Influence of Chloride Ion.-Bray0 observed that chloride ions have a Fig. 12 shows some typical results.

FIG. 12.-The influence of chloride ions on the course of reaction a t 6ooC. [KIO,] : 0-100 M ; [H,O.J : 0.226 M ; [H,S04] : 0.0876 N.

[Cl-] : a = 0.0024 M ; b = 0-0004 M ; c = 0-0003 M: d = O*OOOI M; e = no added chloride ions : f = O.OOIO M ; reaction mixture in

contact with a layer of carbon tetrachloride.

Discussion The Mechanism of the Reaction Between Iodic Acid and Hydrogen

Peroxide.-It has been suggested in recent years that the large majority of chemical reactions take place by a series of bimolecular stages, each of which requires a relatively simple change of pattern of the reacting species.

In the present investigation, the reduction process which leads to the formation of oxygen is governed by the overall equation

5 H202 + 2H+ + 2 103-+-12 + 5 0, + 6H,O. If the above principle is adopted, the deprivation of the iodate ion of its combined oxygen must be visualized as occurring in stages, which involve the intermediate formation of less highly oxidized species :

10,- -+ (10,) + (10) + . Similarly, the oxygen evolved as gas, must, according to the above equa- tion, arise partly from hydrogen peroxide and partly from iodate ions or one of the species derived from them.

It is interesting to enquire how far the kinetics of the reaction studied may be explained in terms of relatively simple bimolecular stages. In an

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Page 10: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARU A N D C. F. CULLIS acid solution of potassium iodate and hydrogen peroxide, there are three species which appear t o take part in the reaction, namely, H+, 10,- (or HIO,), and H,O,. The initial step probably involves either

(a) H,O, and 10,- or (b) H+ and 10,-.

62 5

(a) Interaction between hydrogen peroxide and iodate would pre- sumably take place according to the equation

H20, + 10,- + H,O + 10,- + 0,. The iodite ion may perhaps exist as a transient species but no scheme based on this initial step will satisfy both the stoichiometry of the reaction and the experimentally determined over-all rate equation

(d[OJ /dt) average 0~ m+] ZIIOa-][HzOJ. (b) The second initiating reaction proposed must now be considered.

Iodic acid is not a strong acid ( K , = 0.1686 (25OC) la) and there is evidence that an appreciable proportion of undissociated acid molecules exist in aqueous solution. It is not improbable that iodic acid, like hypo- iodous acid, can in acid solution give rise to a positive ion : 102+ (cf. NO,+, which exists in concentrated solutions of nitric acid).

H+ + 10,- + HIO, + 102+ + OH-. If such a species is formed, it may be assumed that it will react with hydrogen peroxide as follows :

a similar reaction involving 10+ being repeated. would then become

Ha02 + IOa+ + H,O + 0 2 + IOf, The reaction mechanism

Ii z(H+ + 10,- + HIO, + 10,' + OH-).

k2

k3

k4

k5

2(H2O, + IO,+ -+ H2O + 0, + IO+).

~ ( H 2 0 2 + IO+--+ H2O + 0, + I+),

H,O, + I+ -+ zH+ + 0, + I-, z ( H + + OH- -+ H20) ,

which, on addition, steady-state method

[IO,+] r=

[IO+j =

and [I+j =

NOW, 5d[O,]/dt =

whence the relation,

k6 I+ + I- -+ I, gives the required stoichiometric change. If the

is now applied to this mechanism, it is found that K[H '3 [ I 0 -3 - Zi [ FI + J [ I 0 -j -

[OH-I Kw I

z k ,[H 202] [102+ J + z h 3[ H ,O,] [I 0 71 + k H .O J [I +]

is obtained, agreeing with that derived from the experimentally deter- mined average rates of reaction (Fig. 5 ) .

la Fuoss and Kraus, J . A mer. Chem. SOC. , 1933, 55, 4 j b .

2;*

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Page 11: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

626 PERIODIC CHEMICAL REACTION

reaction, i .e.

The reverse oxidation of iodine to iodate

The treatment given only applies to the reduction stage of the overall

zHI03 + 5H,O, --+ I, + 5 0 , + 6H,O.

1 2 f 5H2Oz +zHIO3 f 4H,O does not, however, directly involve the liberation of oxygen, the latter being the criterion that has been used as the basis of the rate equation given.

It has already been pointed out that iodine, which is formed in the reduction process, has a marked influence in promoting the rate of oxygen evolution (Fig. IO), and it will be seen that the mechanism suggested makes no provision for the production of oxygen by this means. In addition to the formation of oxygen as a result of the direct interaction of iodic acid and hydrogen peroxide, it is therefore necessary to postulate an additional reaction, which may be formally designated as

k , I, -+ catalytic production of oxygen.

Iodine is produced by the reaction A6

I+ + I- -+ I,, 41 1 k K + 2 = k6[I+][I-] = "-[H+]2[I0,-][H,0,].

dt Kw so that

On the other hand, iodine is removed from the system both chemically by oxidation to iodate, and physically as vapour by oxygen evolved from the solution. Since the iodine concentration at no stage exceeds z 04 of the theoretical value corresponding to completion of the reduction re- action, it is clear that the physical process of iodine removal (cf. Fig. 3) operates to an extent which is small compared with the chemical process. The maintenance of a low iodine concentration is therefore primarily due to the reverse oxidation of iodine, a reaction which has been shown to occur a t a rate directly proportional to the iodine concentration only.ls Processes leading to the removal of iodine may be expressed as

k* I, -4 removal from system,

where k , is the velocity constant for the average rate of such processes, the value of which will be governed mainly by the rate of the oxidation reaction. Thus it is possible to write as a reasonable approximation,

Since the iodine concentration varies periodically with time (Fig. 3) it is clear that a t any particular instant the rates of formation and destruction are not generally equal. It is nevertheless the balance between these two competing processes which results in the maintenance of a certain average iodine concentration in the system. The value of the latter may thus be obtained by equating + d[I,]/dt and - (d[Iz]/dt)aVerage whence

- (d[IJ/dt)average = ks[Is].

The mode of iodine catalysis is discussed more fully in a later section, but if it is assumed that the catalytic production of oxygen occurs at a rate directly proportional to the iodine concentration,

(d[OJ/dt)catalytic = k,[I2],

l3 Liebhafsky, J . Amer. Chem. SOC., 1931, 53, 2074; see also Bray and Caulkins, J . Amer. Chem. SOC.. 1931, 53, 44.

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Page 12: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARD AND C . F. CULLIS 627 and the rate of oxygen evolution by this means should vary with time in a similar fashion to the iodine concentration. The overall average rate of production of oxygen will however be given by

catalytic

It thus appears that the average rate of oxygen evolution should be governed by the relation

d[021 a [H+]2[I0,-][H202], at both when the reduction process alone is occurring,* and under conditions where iodine catalysis is also taking place.

The mechan.ism suggested therefore predicts a rate law of the same form as that determined experimentally, though in view of the uncertainty of the existence of some of the oxy-iodine species postulated, the purely kinetic agreement obtained cannot be regarded as conclusive evidence for the mechanism.

The Periodic Nature of the Reaction.-In some periodic reactions, the periodicity can be explained in terms of purely physical fzctors, such as supersaturation, the formation and destruction of surface films and the existence of non-isothermal conditions. In other cases the periodicity may be the outcome of the somewhat unusual kinetic features of the reaction, and to this extent explicable in terms of chemical ideas. In this connection it may be observed that there is no thermodynamic condition preventing oscillation a bout a steady-state, provided reaction takes place in the direction of a fall in free energy throughout the oscillation.16

The theoretical possibility of a periodic chemical reaction was first pointed out by Lotka,16 who showed that periodicity could be a character- istic of coupled autocatalytic reactions. Moore l7 extended Lotka’s ideas to a chain of successive autocatalytic reactions in which the concentration of the precursor to the chain remains virtually constant. It was shown that if a non-zero steady-state exists, the concentrations of the inter- mediates do not settle down to their steady-state values but instead oscil- late about them with a slight degree of damping depending on how far from constant is the concentration of the precursor. In such cases, however, the oscillations are sinusoidal in character unlike those shown in Fig. 3.

made a detailed study of the reaction between hydrogen peroxide and iodic acid, and found a somewhat different rate law from that given above. Liebhafsky’s experiments were carried out under conditions designed to minimize the extent of iodine catalysis, i.e. the reacting solutions were kept in contact with a layer of carbon tetrachloride to extract the iodine as fast as it was formed (cf. Fig. 10 and 12). Using Liebhafsky’s values of the rate constants and activation energies for the production and removal of iodine, we have cal- culated the theoretical steady-state concentration of iodine a t 60° C for the re- actant concentrations given in Fig. 3. We find, however, that this value is a t least IOO times lower than the experimentally determined average value shown in Fig. 3. It is evident therefore that continuous removal of iodine from solution profoundly alters the characteristics of the iodate-hydrogen peroxide reaction, so that i t is not surprising that a somewhat different rate law holds under such conditions.

* Liebhafsky

l4 Liebhafsky, J . Amer. Chem. SOC., 1931, 53, 896. l5 Denbigh, Hicks and Page, Trans. Faraday SOC., 1948, 4, 489. l8 Lotka, J . Physic. Chem., 1910, 14, 271 ; J . Amer. Chem. SOG., 1920, 42,

17 Moore, Trans. Faraday SOC., 1949, 45, 1098. 1595.

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Page 13: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

6 2 8 PERIODIC CHEMICAL REACTION One of the objects of the present study of the iodate + hydrogen

peroxide reaction has been to see how far the phenomenon of the periodicity is explicable in terms of chemical ideas, and to what extent physical factors are responsible for the observed behaviour.

It has been shown experimentally (see Pig. 10) that the higher the iodine concentration, the greater are both the rate of oxygen evolution and the frequency of periodicity. In other words, the reaction between hydro- gen peroxide and iodate is autocatalytic, the product iodine catalyzing the reaction which would accelerate to a limit controlled by the purely chemical steady-state concentration of iodine were i t not for the physically interdependent behaviour of the two products. As the reaction proceeds, the concentration of iodine builds up and so the rate of oxygen evolution increases. When the latter has reached a sufficiently high value, the iodine is for a short time physically removed from the solution by the oxygen faster than it is formed. The iodine concentration therefore decreases and the rate of oxygen evolution also diminishes. More iodine is now produced by reaction again resulting in an increased rate of oxygen evolution. The catalytic influence of the iodine produced chemically may be reinforced by that of iodine vapour diffusing back into solution from the space above the reaction mixture. The process can now repeat itself and so give rise to the periodic onsets characteristic of the reaction. I t will be seen (e.g. Fig. 4) that periodic evolution of oxygen only takes place between fairly narrow limits of reagent concentrations. One of the main factors which determine whether or not the reaction is periodic is believed to be the rate of production of iodine ; possible reasons for the disappearance of the periodicity a t low and high rates of iodine formation are discussed in the following section.

On the views expressed above the periodicity is dependent on two distinct factors, one chemical and the other physical. The chemical factor is the autocatalytic nature of the reaction leading to oxygen evolu- tion, whereas the physical factor is the removal of iodine, the product primarily responsible for cata.lysis, by the other product, oxygen.

The possibility must not, however, be overlooked that the periodic removal of iodine is due to chemical rather than physical causes. Bray 9

suggested that, as both the oxidation and reduction reactions involved were autocatalytic, the system fulfilled one of the conditions postulated by Lotka,l* namely that for a periodic variation in rate there must be two coupled autocatalytic reactions taking place. Liebhafsky, who made a detailed study of the separate oxidation and reduction l4

found that the oxidation of iodine to iodate was rapid compared with the reverse process, but that the former reaction was preceded by an appreci- able and highly variable induction period. Chemical removal of iodine is thus rendered unlikely 2s a cause of the periodicity, since in this case, the “ periods ” would be unlikely to conform to the highly regular pattern encountered experimentally. Apart from this, the physical explanation of the removal of iodine as vapour is more consistent with the experimental finding of the decrease in total iodine content of the reaction mixture taking place during reaction (Table I), and with the observed variation of iodine concentration in the vapour phase.

The Mode of Iodine Catalysis.-The results given in Fig. 10 and 11 show that although iodine by itself is almost without effect on the decom- position of hydrogen peroxide, in presence of iodate it acts as a powerful catalyst for this reaction. It is clear that iodine exerts its catalytic in- fluence by increasing the concentration of some species which reacts directly with hydrogen peroxide to liberate oxygen. In the reaction investigated in this work, iodine might owe its action to intermediates formed in the course of its reverse oxidation to iodate, but if this were the case it is not clear why its catalytic activity should not be equally pro- nounced in absence of iodate.

On the other hand, it has been shown by using isotopically distinguished

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Page 14: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

M. G. PEARD AND C. F. CULLIS iodine .I* that an exchange occurs between iodine and iodate in strongly acid solution. This fact is readily explicable if i t is assumed that hydrolysis of iodine leads to the production of iodide and hypo-iodous acid, and that the former then reacts with iodate to give iodine, and the latter dispro- portionates to regenerate iodate and further iodide. In the present reaction, the catalytic influence of iodine may therefore be attributable to one or other of the products of iodine hydrolysis, the extent of which is increased by the presence of iodate.

Iodine hydrolyses according to the equation I, + H,O + H I 0 + H+ + I-.

Either hypo-iodous acid (or its ions) or iodide ions may therefore be re- sponsible for the observed catalysis.

(a) HYPO-IODOUS AcID.-The catalytic decomposition of neutral solutions of hydrogen peroxide is well known to occur through the action of the ion 10-. Hypo-iodous acid is, however, very weak ( K a = 4.5 x 1 0 - l ~

(25' C) IS), so that, since [10-]/[HI01 = K,/[H+], in a solution of pH I, the concentration of hypo-iodite ions may be regarded as negligible. On the other hand, the basic dissociation constant of H I 0 is 3-2 x 1 0 - l ~ (25' C),ao and since [I+]/[HIO] = Kb . [H+]/K,, HI0 should exist in a fairly strongly acid solution mainly as I+, According to the mechanism previously postulated for the iodate + hydrogen peroxide reaction, this species reacts with hydrogen peroxide liberating oxygen, and may thus be ex- pected to promote its decomposition.

( b ) IODIDE 1oNs.-The influence of iodide ions on the reaction cannot be determined directly, since these combine rapidly with iodate ions to give iodine.

Chloride ions have been shown to exert a marked accelerating effect on the iodate + hydrogen peroxide reaction (Fig. 12), the rate of reaction being approximately proportional to CC1-32. Although the mechanism of the catalysis is not known, it is not unlikely that iodide ions would (if their catalytic influence could be studied separately) be found to act in an analogous manner. It may be mentioned that a recent study of the iodate + iodide reaction 2 1 has shown that, in order to explain the kinetic relationships encountered, i t is necessary to assume that the reaction proceeds through the formation of intermediates, IO,+, I0+ , etc. Ac- cording to the mechanism proposed earlier for the iodate + hydrogen peroxide reaction, it is precisely these intermediates which react with hydrogen peroxide to liberate oxygen. The iodide concentration may therefore be a factor of some importance in connection with the catalytic influence of iodine.

I n the previous section it was pointed out that periodic evolution of oxygen only occurs when the iodine concentration lies between fairly narrow limits. At high acid concentration, for example, i t appears that the rate of production of iodine is sufficient to cause such rapid catalysis that the numerous " periods " are merged into an apparently continuous evolution of oxygen (Fig. 4, curve a). On the other hand, at low acid concentrations, the rate of formation of iodine is small. It will be seen (Fig. 4, curve c ) that, under these conditions, the " periods " are separated by comparatively long intervals during which oxygen is evolved at a n apparently steady rate. It has been suggested that the catalytic activity of iodine is in part determined by the concentration of species such as I+ and I-. The concentrations of the latter are normally determined by their relative rates of formation and removal in the course of the iodate- hydrogen peroxide reaction. When their ' ' steady-state " concentration has been reached, a regular evolution of oxygen should take place, but for

Hull, Shiflet and Lind, J . Amev. Chem. Sor., 1936, 58, 535. l9 Skrabal, B e y . B, 1942, 75, 1570. 20Murray, J . Chem. SOL., 1925, 882. *l Cullis, Morgan and Peard, i t i d . , 1951 (in press).

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Page 15: A periodic chemical reaction. The reaction between hydrogen peroxide and iodic acid

630 DEHYDRATION OF HYDRATES the fact that such species are also produced by secondary reactions of the product iodine :

I, + H,O + HIO(I+) + I- + H+ 31, + 3H,O + 10,- + 51- + 6H+.8,

At low iodine concentrations, however, the concentrations of I+ and I- arising from the above equilibria will be less than the normal “steady- state ” concentrations, and in these circumstances, oxygen is evolved at a steady rate. It is only when the iodine concentration exceeds a certain value that the concentration of the related ions begins to rise. When this occurs, the rate of oxygen evolution starts to increase, and a ‘‘ period” eventually takes place. At even lower acid concentrations (Fig. 4, curve f), the iodine concentrations never rise sufficiently to enable the active species to exert their catalytic influence, so that a steady evolution of oxygen persists throughout the reaction.

In conclusion, therefore, it seems that the mode of iodine catalysis cannot be established with certainty. It is probable, however, that the action of iodine is not ascribable to any single entity but to a system in which several inter-related species are involved. It does not seem possible a t this stage to produce even a simple mathematical treatment of the periodicity encountered in this reaction. There is no doubt that it is a genuine and reasonably reproducible phenomenon, but it occurs in a system in which a very unusual combination of chemical and physical factors are operating. In view of this, it is not surprising that so few reactions of this spectacular type are known.

Physical Chemistry Laboratory, Oxford.

23 Luther and Sammet, 2. EZektrochem., 1905, 11, 293.

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