The kinetics of the addition ofhalogens to unsaturated compounds

Item Type text; Thesis-Reproduction (electronic)

Authors Bryan, Elmer Leo, 1900-

Publisher The University of Arizona.

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THE KiraiCB OF THE ADDITION OF HALOGENS TO DESATORATED COMPOUNDS

ty

Elmer L, Ir^n

A The BisBubsaitted %9 the faculty of the

Department of Chemistry' - ' ' ' ' - ' ; ' * ' .V " ■ >

in partial fulfillment of the requirements"far the Aegree of

Master of Scteaae

in the Graduate College University of Arlaona

1938

Approved:Major JProfewte 4 ^

: to #####& $D mmow&s at

#9#

«St*8 *$ rssu-'

X

• 2-

ACOOl/LEDQ&EHT

The writer wishes to express hie 6iost sincere gratitude for the generous advice and assistance of Dr. Lathrop Emerson Roberts, under whose direction this investigation was made.

1 4t.V02

TJIELE OF COmUVxD

FagaIntrofluetioB 1Review of the Literature 0General Experimental Procedure ISPreparation of Apparatus and Materials 17

Apparatus 17Preparation of Reagents 17Preparation of Eolventc 18

Experimental 20Blseutsion of the Results SOEuncmry 45

46Bibliography

Introfluetlon

From tho study of the rates of chemical reactions and the resultant mathematical relations the mechanism of reactions can often he determined* From such studies general principles can often he formulated and expressed mathematically, as for example the Arrhenius equation (In theequation of Lercis ( , and others. The eol- lieion theory of reaction velocity is an important development of kinetic studies* From this theory Lewie1 derived bis equation for calculating the value of the reaction velocity constant. The collision theory of activation is now accepted as the host explanation of the kinetics of gaseous reactions* Such studies are concerned with the fundamental nature of chemical change and offer an interesting field for investigation. Although much work in this field has been reported, many questions remain unknown and considerable uncertainty still attends much of the theoretical treatment of the problems involved. Especially Is this true with reactions In the liquid phase where almost no general principles have been discovered,2

The first chemical reaction to be studied from the kinetic viewpoint was the inversion of cone sugar in aqueous acid eolation.3 Since this work numerous papers have been published in an effort to solve the kinetics of the reaction. Its

precise mechanism Is still unknown, a fact which gives some iflea of the difficulties attending such Investigations. Progress in chemical kinetics from the time of the first work wao clow; in fact, for a half century little was accomplished. This was particularly true for reactions in solution where many factors complicate the problem, such as formation of chemical complexes^ reaction of solute with solvent,5 ionization,4 uncertainties as to the correct application of theories of activation and deactivation,4 and of redistribution of the energy of activation among the internal degrees of freedom,4 propagation of reaction chains and side chain reactions, motivation by means of radiation,” peroxide effect of unsaturatod bonds,6 the effect of surface tension,7 catalysis of solvent and impurities,8 and others, perhaps some of thee yet undiscovered or described. As the influence of these complicating factors was studied by various investigators progress became more rapid. Host of the difficulties mentioned, however, do not enter Into the study of gaseous reactions. It is due to the greater simplicity of reactions in the gas phase that investigators of kinetics became particularly interested in gaseous reactions.8 Even with gases investigation is often difficult. Thus the gases studied do not obey the ideal gas laws; the taehnique of handling them is difficult; the rate and mechanism say be influenced by the walls of the reaction chamber, by the initial pressure, by the type of material composing the reaction

8

ebmm&er, by impurities and other factors; there may be difficulty of determining the identity of reaction products; the complication of opposed reactions, side reactions and chain reactions nay be present.

the development of the collision theory of activation is one of the most important contributions to chemical kinetics. Since its development the advancement of knowledge of chemical changes has been made more rapid, this theory is now accepted as the most satisfactory explanation of the velocity of gaseous reactions# The theory itself is the result of a succession of developments. Arrhenius first stated the relation between the velocity constant and the absolute temperature: In k=B-^. Several other empirical equations have been advanced but the Arrhenius equation is the most

Jf/fTT/ eatlsfaetary. The*equation k»Zc- , where Z is termed the critical increment of energy and 2 is the collision number, was derived by Lewie from the application of the I'azwell- Boltzoann distribution law. lewis showed the hypotheses to be substantially correct by actually calculating from theory a value for the specific velocity constant which agreed with that obtained from experimental data.1 The ideas of Perrin,® who believed that a unimolecular process was one whose rate was independent of the pressure of the reactant and that isolated molecules should decompose at the same rate me whenpresent in a group of molecules, led Lewis10 to postulate that the reaction velocity was affected by radiant energy.

4

Both of these hypotheses have been discarded. Llnaemanft,1*In answer to Perrin’s hypothesis, assumed that molecule® &e- ocMBpeelng unleolecularly could still be activated by collision If the activated molecules existed for a finite time before most of them reverted to a normal state while the rest decomposed, i’hi® theory involved two assumptions: (1) that at low pressures when there are so few molecules for the space occupied that the activating collisions fall off ;the uniraolecular constant should fall off; {2} that at high pressures the rate of activation should exceed that for a uni- moleeular reaction. Hlnehelwood*3 ales showed that simple molecules should be expected to decompose blmolocularly and more complex molecules to decompose unleolecularly* which coincides with activation theories.

She application of the collision theory of activation has aided greatly In the investigation of the kinetics of gaseous reactions and lead to renewed Investigation of reactions in the liquid phase. She collision theory had to be developed for reactions in the liquid phase. In gaseous reactions activation may be accomplished by collision of reactant molecules or by collision of reactant molecules with the walls of the reaction chamber, in solutions in the liquid phase collision of solute molecules with solvent molecules with solvent molecules must also be considered.'fhe expression l2o = | ^ ^ 7gt!r) has teen developed by Jowott14 for the number of these collisions.

8

Recently much study has been given to the kinetics of reactions is solutions, these Investigations have shown that there is no fundamental difference between reactions of the liquid phase and those of the gaseous phase,3,6 In all reactions which have been earrlefl out in the two systems this conclusion has been verified in those cases where the solvent was an inert, normal liquid. In general, however, the solvent has been found to exert a positive or negative catalytic effect. Comparison of such reactions in the liquid phase with the same reactions in the gaseous phase has shown the degree to which the solvent has affected the reaction, The study of the causes and mechanisms of the effects which solvents exert is the problem of contemporary investigators in this field, She mechanism of a reaction In the liquid system is usually difficult and frequently impossible to determine. As yet very few general principles have been discovered, Most of the general statements that have been made are misleading for they are drawn from specific reactions, Ibis is shown by the fact that the catalytic order of a series of solvents may differ evon for similar reactions.16

A study of kinetics usually Involves the attempt to determine the reaction velocity constant. % l s constant is calculated from various equations, characteristic of the order of the reaction, the degree of completion of the reaction and other factors. For a complete reaction In general.

nhere nsn, on© rmeli equation is kCa-xKb-slCc-x) - - -f is the colee per liter reacting in time ntn, end nan, "b",Mc % etc., represent the initial concentrations of the reactants in moles per liter. For a unitaolecular reaction b*0. o*0# etc.; for a bieoleeular reaction e*0, etc. The integrated form for each order becomes specific. Since, in thte paper, wo are interested only in second order reactions where &A, «e get, k- -(frfxf log This equation will beapplied later In treatment of experimental data#

For opposed reactions of second order reaction a more complicated equation rccults* The opposed reaction between the resultants may be enieoleoular, blmolecular, termolecular, etc*, all depending upon the products. For the reaction chosen for study in this paper we may consider the ease where the opposed reaction, that is, the decomposition of the addition product, is untmoleoular# Here wo apply the differential equation = k^a-x) Cb-x) - IĈ r. Since at equilib-

rlura - 0, it follows that (b-x^) * kg^or kg =^1 - k^Kg, where ICe is the equilibrium constant

xe(alco|oI|(iodine). Substituting this value of k2 in the above.

we get ||- = (a-x) (b-x) - Kex * Integrating we get: •

SSiis equation is used in calculation of constants from data given later. For brevity the reaction velocity constant for

f

a reaction Milch goes to completion will he referred to as ”kn; and the reaction velocity constant, second order, for opposed reactions will be referred to as nkxn«

Calculation of satisfactory velocity constant from the data by the equations given above does not necessarily show the order of a given reaction, The reaction of bromine with hydrogen peroxide is an example.1? Here a parallel reaction of Br“ ions with hydrogen peroxide causes a stage in the reaction to be reached in which the bromine becomes constant. Another illustration is ionic reactions subject to the electrolytic effect.18 The most reliable method of determining the order of a reaction is from the half time. The general equation for the half time is where nn" is theorder of reaction.

Most reactions take place in liquid systems and these are mostly bimoleeular. It is therefore desirable to know sore about the kinetics of such reactions than is now known. Such a study was chosen for the investigation reported in this paper. It was decided to study a reaction which would be reasonably sure to be second order. The halogenstion of the ethylene bond appeared to bo such a reaction. The search for a suitable compound containing the ethylene bond led to the selection of allyl alcohol. The structure of the molecule of this compound seemed to indicate that it would be soluble in a variety of solvents. It was found to be soluble in fifteen different solvents of varying types. The tolling

e

point of allyl alcohol, which is S4,5°C,, makes it convenient to work with as it is not too volatile* She greatest disadvantage of this compound Is Its lachrymatory effect*

9

Rg-gtew of the literature

The klnotics of the addition of halogons to the ethylene bond have been studied by a nmaber of investigators in both tho gaseous and liquid phases* Zost of the investigation, however, has been done in gaseous systems* The effect of

t/ coating the walls of the reaction chamber with parai^n and other substances has been studied by Iforrish20 and Keisig.21 Mltsukuri,22 Williame,23 Kimreuther,24 Stewart and lund25 and others have investigated the reaction of bromine and ethylene in the gaseous phase, A few bromine additions to various mass tura ted compound# have been studied,26*27 in liquid systems, mostly in carbon tetrachloride, carbon bisulfide and chloroform as the solvents* The rate of reaction between bromine and unsaturated aliphatic acids has been studied as evidence of ■terieisomcriem*28*29 Photochemical investigations cf bromine additions to unsaturated compounds have been made.30*31 Activation energies have also been studied by various investigators*27

The addition of iodine to the ethylene bond has been studied very little. Some investigation has been done in the gaseous phase.32 Schumacher believed that he had determined tho mechanism of formation of ethylene iodide in this phase.33 But a search of the literature for Investigation of the addition of iodine to unsaturnted bonds, that of ellyl alcohol

10

in particular, proved almost fruitless* Most studies of this reaction have been done from the organic viewpoint, not from the kinatie* Products of the reaction have been determined by several investigators.

Polisson34 did come work on the kinetics of the addition of iodine to ethylene in carbon tetrachloride solution. However, as stated, most of the investigations have been eomoermed with bromine additions and the products of reaction* B#r% and Mylius35 were the only investigators found who had studied the reaction between Iodine and allyl alcohol in liawid systems. They studied the reaction in only three solvents, carbon tetrachloride, carbon bisulfide and chloroform. By using the blmolecular formula for the velocity eonsteat "k", which was developed in tho introduction, these Investigators concluded that the reaction went virtually to completion and was of the second order. Reference will be made a number of times in this paper to thcce investigator*.

The allyl alcohol used by Hers and Mylius was refluxed three hours over calcium oxide, then once over metallic calcium. The chloroform, carbon tetrachloride and carbon bisulfide solvents were all C. P. and refluxed once over calcium chloride.

During the course of the experiment, samples of the re- ' action were titrated for Iodine determination with sodium thiosulfate, using starch solution as an indicator. After the endpoint was reached they noted that a blue color slowly

ureappesrede

A few results of their experiments are given in Table 1.

TABLE 1

1* Solvent: Chloroform

(«) a*4.866 b-0.6279

t9 Er,z

50 ter*(b) **2.985

0.13120.4780b"0^847

t16 Hr. 0.1H9825 « 0.278931 « 0.296044 " 0.357970 • 0.4691

(c) B=1.0515 b=0.5544

a=niinmol8 of allyl alcohol in 20 c.c.eolotio#.

hsGillimolB of iodine in 20 c.c. of eolation

k x 10-3

2.9

' 5.22.7 2.42.7

(a)

t46 % . 120 n 144 "214 n 312 wa»0.5945

0.20980.35630.38350.40050.4206b-0.5344

I #4.94.93.62.8

38.5 ‘Br. 72 n 102 «

141.6 " 196 w 270 "33® *

.0*450.16330.19300.23110.24500.28310.3078

k % 10-34.04.43.93.83.02.92.8

u

M L B I (Concluded}

2. Solvent: Corkon bisulfide(a) a=4.411 5*0.539

1 a k x 10-3? hr. 0.1911 6.39 n 0.2499 7,016 " 0.3685 7.619 ” 0 ^ 8 2 1 6.9

26.5 " 0.4312 . 6.4m 8*2.481 5*0.539

1- a - k % 10“3- 20 hr. 0.1400 2.8

45.5 ” e.mm# 3.054.5 " 0.3107 2.967.5 R 0.3322 2.7

3. Solvent: Carbon tetrachloride(a) 8*4.375 b=0.5432

i a k„x 10-24.5 hr. 0.1813 1,0

5 0.2322 1.16 » 0.2514 1.10 a 0.8028 1.010 " 0.3332 1.0

General Experimental Procedure. . . - • - . ■ - -• . . .

Preliminary experiments were made to determine which halogens to use* Since the reaction was to he studied in liquid systems fluorine and chlorine were unsuitable* Preliminary experiments with bromine and allyl alcohol showed that bromine was too active, the reaction going to completion so rapidly that it was difficult to follow by removal of sample® at different intervals and analysis. Preliminary experiments with iodine and allyl alcohol in carbon tetrachloride showed the reaction to proceed at a measurable rate. Iodine was also found to be soluble in a number of solvents of different types.

$he general procedure used in studying the kinetics of the reaction of allyl alcohol and iodise in liquid eyctcms Is comparatively simple.

Since the solution is accelerated by light it was carried out in the dark, and at a constant temperature of 25°C* The various solvents used were carbon tetrachloride, chloroform, carbon bisulfide, ethyl alcohol, butyl alcohol, benzene toluene, acetone, dloxene, propylene chloride, iso-propyl ether and tetrachloretbane. The initial concentrations varied from 0.5S* to 2M* 100 c.c* solutions of each reactant were used.

The 100 c.c* of solution of each reactant were added to

14

the reaction flas’.;, quickly shaken anfi replaced in the thermostat, the time of mixing "being noted. Then, as the reaction proceeded, 10 c.e. camples of the reaction mixture were removed at measured intervale "by pipetting and quickly added to prepared eamplc bottles to stop the reaction. In a few of the first experiments the reaction was stopped by placing 10 c.o. of sofllm® thiosulfate of known concentration in each cample bottle before adding the reaction mixture to remove the exceas iodine of the sample* The excess sodium thiosulfate was then back-titrated with iodine colutien, using starch indicator, to determine how much Iodine was used up in the reaction with allyl alcohol. However, In most of the experiment# the reaction was retarded by adding the 10 e.c. sample to 100 c.e. of water. Then the excess iodine @f the sample was titrated directly with sodium thiosulfate, using starch indicator, to determine the amount of iodine used in the reaction with allyl alcohol. From the data thus obtained, *k" or "k%" was calculated for each sample from the velocity constant equation# gives in the introduction.

The equation for the reaction velocity constant was simplified for each reaction so that the number representing the cubic centimeters of the titrating agent could be substituted directly into the equation. The method used is illustrated in the following example.

(2) a=0.10102=inltial concentration of allyl alcohol mols perliter.

15

(5) b»0.0501?»£nitial concentration of loSine in mole per— ---- liter

(4) a*h».05085let v=j/o. c.c* of .09656 1. eoaium thiosulfate used in titration.

(5) Equivalents of excess iodine in 10 c.c. samples.00009656 v Then the sols per liter ef excess iodine in the samples

* .- x 1000(6) = .004828 V = t-x

z » h-.004828 v(7) .05017-.004828 v, subst. from (3).

a-x = 0.10102-(.05017-.004828 v), eubet. from (2) and (7).(8) = .05085 + .004828 f

»' « '' = « ■ « • « -(2), (3), (4), (6), and (8) Into (1).

(10) 85^.004828 T ) simpUfyins (9).

The equations for the specific velocity constant for opposed reactions involving the number of cubic centimeters 2 3 4of the titrating agent arejdeveloped in the following

(i) 4,

(2) a=0.10232* initial concentration of allyl alcohol inmol® per liter.

(3) b=.04997, initial concentration of iodine in mols per‘ liter.

(4) a-b = .05235

16

Let v * no, c.o* of .09126 K* sodium thiosulfate used Intitration- . '

(5) Equivalents of excess iodine in 10 o.o. sample» .00009126 V

(6) Then the mols per liter of excess iodine in the sample *. .. 'R0992126_J[----- X 1000 « .004563 V « b-X

(?) x . b~,004563 v # .4997-.004563 v , oubot. from ( 3 ) .

(8) 2x m .09994-.009126 Va-x » 0.10252 -(.04997-.0045632 v)» subst. from (2)

and (7)(9) a-X « .05235 * .004563 V

(10) Ve « 5/05, no. c.o. of sodium thiosulfate used in titration at equilibrium

y m , where K0 is the equilibrium con-xe stant and x* the mols per liter

reacting in tine t .* subst *

(11) K * ,.06450U(a + b + Z ) m 0.10232 + .04997 » .06460, subst. from

(2), (3), end (11).(12) » 0.21679

V(a + b + K )2 - 4 ab «V(0.21679)B - 4 x 0.10232 X.04M? 0 oubst. from (12), (2), and (3).(13) V(a + A + Kq)2 - 4 ab » 0.16294(14)

kl-mzWT[tolvW^s~+I-tsl72\ tihero l.tS17z to a logarltlm.

17

Preparation of Apparatus anS Materials

The ApparatusThe reaction flasks in these experiments were 300 c,c.

glass stoppered bottles of brown glass, which filtered out most of the light. Since this reaction is photochemical, further precautions were taken to prevent light from entering the reaction flasks by painting them with several coats of opaque, black asphalt paint before use. These bottles were thoroughly cleaned with cleaning solution and dried before using. This was necessary to lessen the probability of catalytic action or surface effect.7 The same precaution was taken with all volumetric measuring flasks and pipete used.

The thermostat was a large water bath, well stirred, electrically heated and automatically controlled. Tcnncra- ture control wan accurate to within .05 degrees.

The apparatus used in distilling all solvents and the allyl alcohol was made entirely of pyrex glass with the exception of a five-inch immersion thermometer* which was of a different glass, and a two-inch piece of platinum wire used te hold the thermometer. All joints were of ground glass.

Preparation of ReagentsThe allyl alcohol used in the reaction studied was

18

obtained from Eantsum Kodak Company. It hafi a boiling range of 95.5° to 97°C» It was refluxed for five hours over Kerch’s Reagent calcium oxide, to remove all moisture. It was then distilled from the lime* retaining that portion distilling over at S4.5°C. Two further fractionations at 94.6°e« were believed sufficient for the purpose of these experiments. It it as feared that moisture might be a catalyst.

The iodine used in the reaction studied was Merck’s Reagent. It was further purified by grinding with Merck’s Reagent potassium Iodide and subliming the iodine. This was done to remove all free chlorine and bromine, which would appreciably affect the velocity constant. % o more sublimations followed in order to remove any possible potassium iodide that might have sublimed with the iodine. The iodine thus purified was stored in a deeaicator charged with calcium chloride.

*The solvents were nil treated in the same general man-

ncr to remove all moisture. The ethyl alcohol was absolutealcohol furnished by United States Industrial Alcohol Company, while the carbon tetrachloride, chloroform, carbon bisulfide, butyl alcohol, benaene, toluene and acetone were all J. T. Bakers Analyzed, C, P. chemicals. The remaining solvents, dioxanc with a melting range of 10.5° to 110C., propylene chloride with a boiling range of 95° to 9B°C.,

iao-propyl ether with a boiling range of 67° to 69°C. and

IS

tetraehlorethttne r/ith a tolling range of 144° to 14G°C., were all eeeureS.from Bastman Xodak Company* Each of thesesolvents was refluxed for four hours over J, T* Eater’sAnalysed C* P. stick calcium chloride, to remove all moisture.and fractionated three tines,^ retaining those port! the following temperatures:

Solvent ■. ; Tcmpe:Carton tetrachloride,........f3.50-74.5°C.Chloroform..................57d-58.55C«Carbon bisulfide.............43.8° -44.8°C.Ethyl alcohol................76°-76.20C.Butyl alcohol.... ...........115°C.Benzene............ .77°-77.50C.Toluene.. ..... 107.5°C.Acetone.... ..........54.5°-59.5*0.Diozane................... ,..S8.5°C.Propylene Chloridc».*....»,»*9^.52c* *Isopropyl other...........,..65.6 -65.8:C,Totrachloretiianc,............ 140 -143.5 C.

at

The boiling temperatures in the above are those actually recorded. The elevation was 2350 feet above sea level with a barometric reading varying around 700 sum*

Blank experiments were made with each solvent, i.e., preparing a known concentration of iodine and titrating samples taken at intervale to detect whether iodine reacts with the solvent. The solvents selected above did not appear to react with iodine.

20

Experimental

fhe data herewith presented are all for the reactions at 25°C. Preliminary experiments were made with the different solvents to determine the suitability of the solvent for a study of the reaction of iodine with allyl alcohol. The solvents all proved suitable with the exception of acetone.The reaction was practically complete in acetone in three minutes. An odor resembling iodoform wan present. These ehsraeteristics mado acetone undesirable and further investigation with It was discsntMusd.

In several preliminary experiments with the same reaction In carbon tetrachloride as the solvent, an induction period was noticed. This varied somewhat according to the initial concentrations of the reactants, the average being about one and one-half hours. This Induction period may have been present in all the reactions in the various solvents, but since the first samples were removed in most cases after an elapsed time of over one hour, except for one experiment in sfueous solution, it was not definitely confirmed in a preliminary experiment.

The presence of this Induction period In a preliminary experiment Is shown In table II, Each 10 c.c. sample removed from the reaction flask was added to 10 c.c. of sodium thiosulphate to stop the reaction. The excess sodium thiosulfate was back titrated with iodine solution as explained

21

above. The reaction goes practically to completion.

SABLE II,SolventCarbon tetrachloride. Concentration of Iodine:— .vonoentration or iodine:— .0264K, per liter Concentration of allyl alcohol:— .0252M. per Aqnooun solution of sodium thiosulfate:— 0.1< Atmeoun solution of i o d i n e 0.10081!.

liter 1032!?

v * c.c, of iodine solution used in titration.t X k x 10’

21 min. 2.7741 w 2.771 hr, 1 n 2.681 " 20 u 2.78

1 " 41 it 2.781 M 56 u 2.84 3.6522 n 18 * - 2.88 3.5812 » 40 * 2.88 #.#1S » 0 # 2.89 2.8324 ° 7 # 2,90 2.0646 n 13 u 2.97 1^197 " 3 * 3.07 2.0198 n 16 n 3.19 M l #10 ” 25 # 3*25 2.08212 « g * 3.3825 K 50 *' 4.20 8.585m " 24 # 4.40 3.04447 » 4 n 4,8871 * m II 5.16

She results in this experiment indicate that the union of iodine with allyl alcohol in carbon tetrachloride, after an induction period, proceed slowly by a Mooleculsr reaction.

An shewn in Table II, the reaction in carbon tetrachloride was slow. The attempt was made to speed up the reaction by means of a catalyst, Mercuric chloride Is used as a carrier of iodine in fats and oil36 and was therefore first tried, Mercuric chloride, however, was * appreciably

22

soluble In carbon tetrachloride hat was soluble in absolute ethyl alcohol. $hie led to several experiments with mercuric chloride for the reaction, the reaction was greatly epeefled up, the rate being almost proportional to the amount of catalyst present. Upon calculation of the velocity constant for a reaction proceeding without an opposing it was found to drift downward fairly rapidly aa shown in table III. this did not agree with the results of Hers and Mylluc36 given in Table I, although they did not use absolute ethyl alcohol as a solvent. The reaction was stopped by adding the samples to sodium thiosulfate.

T&BLE IIISolvent&bnolute ethyl alcohol.Concentration of allyl alcohol:— 0.0526%.Concentration of iodine:— .0504%.Catalyst:--!.25 ga,of mercuric chloride In 100 c.c. of

reaction mixture.

1 z k x 10"*65 min. 34 8@e. 2.33 16.0211 # 40 2.S9If n 36 n 3.31 10.1022 *. 81 n 3.4629 if n 3.68 5.96m # 40 3.8047 n tf it 3.S6 3.761 hr. 0 * 23 it 4.091 " 16 n 29 4.27 2.34

1 n 32 24 4.341 r 52 n 53 n 4.46 2.382 " 17 # f it 4.63S ” 16 it 80 4.73 1.364 11 43 15 4.826 n 19 « 41 4.92 0.7067 " 56 e 55 ft 6.0110 » 0 n 13 6^)9 0.457Final 5.20

ante, solvent or catalyst eight have eausefi the constant to drift, numerous repetitions of the reaction was made with varying purifications of the reactants, solvent, and catalyst with the same result. These tables show that the reaction did not go to completion, and that mercuric Iodide and cadmium chloride have only a slight catalytic action.Ho constants were calculated for these reactions. A few times a brick-red precipitate appeared at the endpoint of titration which disappeared upon adding additional sodium thiosulfate. This precipitate was believed to be mercuric iodide.

At this stage two other catalysts were tried. One particular characteristic of mercuric chloride which differentiated It from most inorganic compounds was its slight Ionisation. It was thought that other compounds with this characteristic might be more efficient. Two such compounds are mercuric iodide and cadmium chloride. These two catalysts proved lesE efficient than mercuric chloride. This may show’ why no constant for a complete reaction was obtained* The reaction did not go to completion in absolute ethyl alcohol with these two catalysts, although, as shown later, the equilibrium point was shifted nearer to comple-

/ tion. This Is contradictory to the results secured by Hers and Bylius as shown In Table I. As stated above, they had not need absolute ethyl alcohol as a solvent.

fhe results of the reaction of iodine tilth allyl alcohol in absolute ethyl alcohol as the solvent and tilth mercuric Iodide and cadmium chloride as catalysts are shown in Tables IV and V.

CABLE IVSolvents— Absolute ethyl alcohol Concentration of allyl alcohols-*OS3SU*Concentration of iodines— *0S22E.Catalysts— 0.4258 go. of mercuric iodide in 100 e.e*

of the reaction mixture.te reaction was stopped

sodium thiosulfateby adding the sample to

i x3 min. 0.929 « 1.0619 • 1.1240 " 1.151 hr# 54 ” 1.36

4 17 ” 1.696 n 11 * 1.8210 is 9 " 2.2121 * 47 ” 2.6446 # 26 " 2*71 ' •70 * 2 " 2.79194 0 R 2.81For the completed reaction v » 5.20*

TABLE VSolvents— Absolute ethyl alcohol.Concentration of allyl alcohols— .09S8M.Concentration of iodine.0497H.Catalysts— 0.488? gra. of cadmium chloride in 200 e.e. of

the reaction mixtureThe reaction was stopped by adding the sample to 100 c.c.

of water.

1 hr.2 "

t1 min. 7 ©in. 22 n

48 »17 "13 "

9.95!:!i6.968.438.00

*

*

25

TABLE V (Concluded)

i 14 hr. 10 min. 7.086 " 47 " 6.5011 " 21 * 5.6125 " 2* « 4.6936 * 44 " 4.5147 ” 62 " 4.4071 " 5 n 4.88u * 10 ” 4.115 days 4.16For a complete reaction v = 0.

Mother trial reaction with ethyl alcohol as the solvent and mercuric chloride as the catalyst was made to tcct whether the reaction reached an equilibrium and* If so* whether a second order reaction velocity constant for opposed reactions could be obtained# Ac shown In Table VI, no satisfactory constant was obtained. This sac believed duo to other complicating reactions.

TABLE VI.SolventAbsolute ethyl alcohol.Concentration of allyl alcohol:— 0.1005%.Concentration of iodine:— .0*17?%.Catalyst:— 0.9638 gn. of mercuric chloride In 200 c.c.

of the reaction mixture.The reaction was stopped by adding the cample to 100 c.c.

of water.t V

2 %aln. 4.597 « 4.88 0.6099

17 " 8.T8si : 3.222.67 0.21041 hr. 12 " ■ 8.8V ■ 0.11771 " 58 n l.Sf

4 n 25 “ 1.03 .05502

26

CABLE VI (Concluded)

t * %7 hrv 5 tain 9 " 47 n

22 ° 35 1145 »

0,77.03313

0.330.220.130.12

For the completed reaction v - p*

The 4ieco7ery that the reaction of iodine with allyl alcohol In ethyl alcohol did not go to completion led to the trial of other solvents of various types. As stated above, while employing mercuric chloride as a catalyst a brick-red suspension had been noticed at the endpoint of titration of sampleo. This had happened in those instances where 100 c.c. •of water had been employed to stop the reaction. This brick- red coloring interfered with detecting the exact endpoint.It was further feared that complicating reactions were resulting in the formation of the brick-red precipitate, all of which made the presence of the catalyst objectionable. Consequently all further experiments were made with no catalyst present, other than the possible catalytic effect of the solvent itself. This must be remembered in interpreting the following tables.

At least throe experiments were performed with each solvent at different concentrations for the reactants. It was deemed beet for the reaction of the samples to be retarded by

2?

®aaitlon to 100 e«e. of water, $hie unually aided in determining the endpoint, therefore In the following the re- aetlon stepped by thlo Beans,

tables VII to X?X$ present the reemlte of experiments on velocity of addition of iodine to allyl alcohol In various aolventa In the. absenee of any eatalyet. In those experiment® in which the reaction went to completion or nearly so (l.e,, numbers VIII, XX, X, XI, and XII) the constant *k* for the second order.unopposed resetIon is calculated. In those experiments In which the reaction reached an equilibrium as shown by the constant value for the volume of sodium thiosulfate used (l.e., number# VII, XIII, XV, and XVII) the constant for the oppoaed reactions Is shown. In seew easesboth constants are given for comparison. In calculating the constants in tables VII to XVII, the following formulas, discussed on page';six, were used, The general forme are:

> 1 ewhere A, B, € and D are constants specific for each reaction, *iw la the time in minutes and "v” is the number of cubic centimeters of sodium thiosulfate used In titration. The values for these constants and variables are given in each table. The Initial concentrations in taols per liter of reaction mixture arc given In the tables for iodine and allyl alcohol. These are noted by na* and "b*.

HO

gggsgup

•TABLE .VII

Solvent: Atwlmt#a s 0*1030%. allyl alcohol A = 0.1007 B = 33.90 C

1 hr. 10 min.56 8

101

2717648355500

fcl x 10-3

6.715.765.67

I I5.61

6.7486.5656.448

5.331

520%. allyl alcohol 10 B s 16.58

t 1

•0499M. loainc B « 9.18617-10

x 10-5~ o Din. 8.370 ” 7.96 » 6.2600 n 7.33 6.2350 R 6.66 5.61734 8 6.40 5.2890 " 6.36 5.6510 n 30 «

6.286.24 5.184

50 83

6.216.20

5»BIB m iSolvent: Carbon tetrachloride

a: 0.1010%. allyl A; 45.25 E: .05085

124610122326313647

hr .""41 ” 50w 46« to" to" 51” 41" 40n 26 * 24" 0

aim. ■$C4$8.91

m1.9?1.46

B:.009722

I10.339.571

a;As 1811 B: .00127 C:

3t

hr.""d min. 7̂6 n W * 7.118 ” 30 " 6.S313 « 7 11 6.6026 " 21 " 5.38: to w 40 " 4.93to " 40 “ 4.7736 11 30 « 4.4948 n 30 w 3.3956 ,r 45 " 3.0073 " to " 2.38126 • 2 ” le48

vas2il1.1831.167i:SS1.3051.3281.287

5ABL2 IXSolvents Chloroform

1* a:A; 47.5 3: .050 C: .004563

t V5 hr. 16 min 10.0810 ” 45 " 9.5324 * 55 “ 7.4030 " 52 "47 « 33 ” 5.3085 ” 37 51 4.8573 » 56 » 4.0384 " 25 » 3.7596 « 16 » 3.50

•049SM. Iodine B:

k % 10"5

2*7162,6432,6415.0523.0743.0132.759

2. as .051511. allyl alcohol AS 1437.8 B: .0016

12 hr. 24 min. 9.1028 » 27 " 6.9288 " 40 ” 6.2155 « 4 » 8.0472 .« 32 7 4.4698 " 47 " 8 9 5118 w Final

45 * 3.8#2.89

h 049911. iodina C: .004563 B: .0047093r k x 10~3

5.2366.5606.8116.7446.3295.8115.419

31

tTABLlT X

Solvent: Carbon M e u i n g es: O.IOOB!. nllyl alcohol For k: A; 45,881 B: ,05013For kx: A: .07772 B: 15.25

b: .04991?. ioQlne C: .004563 B; .0091454 C; 1.78 B; 9.54472

5 hr11 "

t v k % IQ"3....... .

47 Bin. 9.27 4.96259 * 7.87 4.9476 * 3.54 . 4.8761 # 4.92 4.80935 * 3.94 4.35249 # 3.42 4.12429 # 2.83 ■ 3.5001.78

kx a 10-5

2.2162.2212.2752.2902.1782.2171.882

.0807911. ally! alcohol k: A; 2190.5 B; .00105 b: .0S573M. iodine C: .004563 B: .0046593t 5 k % 10*3

hr.

Final

30 Bin. 10.65 2.92824 ” 10.20 2.*##U n 9.28 2.36229 « 8.75 2.26410 " 7.85 2.10645 B 7.34 1.89241 ■ 6.69 1.70024 ■ 6.11

5.341.604

fABIS XI

U s: 0.10024S, allyl alcohol For k: A: 45.635 B: .0504

Solvent: Benzene

: t12 hr. 0 min.

V

5.9635 » 56 « 9.0060 « 24 tl 8.2684 " 10 ft 7.69108 R 33 ts 7.26145 ” 0 It 6.51154 “ 15 u 5.79230 " 29 # 4.80Final 4.10

b:C:

•045821!. Iodine .004563 2>s .005181

7.510

! i5.630

2. a: .08496*.For ks 5; 8S .02512

t V23 hr.~l2 min. 1(C8854 ” 10 " 10.8783 " 51 " 9.99121 » 3 ” 9.51168 " 15 " 9.18

217 ” 16 ” 8.78311 n 23 « 8.16Final 7.20

k^z IQ-*

8.7628.2579.014

9.078

TABEE XXI

1. as .09975M. ally! aleobol For k: As 45.845 B; .05017

Solvent; Toluene

11 Hr.55 53 72 96 155 216 423 Final

2 min. 22 40 22 0 25 10 14

10.158+697+536+775.994.603+732.301.90

b:*

iii i i1.0340.808

a: +02414M. allyl alcohol k = 90.302,

b: .0496m. iodine

t V24 h r ” 9 min. M n 41 *, %145 « 2 »192 n 46 8*89240 « 42 « 8.50312 « 58 8.07431 n 27 ” 7.49M m l ?.20

k X 10-3

0.9100.9470.9520.980

34

TABI.E H U

Solvent: Butyl alcohola: 0.10232%. allyl alcohol A; 0.16254 B; 30.84 C; h;5.05

1 hr* 21 min. 8.466 n 58 « 6*168 ” 11 « 5.5811 « 18 " 3*3124 " 32 ■ 5.0930 * 63 » - 4.9848 '* 35 « 4.9#77 * 0 " 5.11

Bs 5.15172 kj x 10-316.14 5.955 5.062 9.496 6.775

2. a: .0457511. allyl alcohol hi. A: 23.72 B: 14.574 Cj 6.30

t ■ T

2 hr. 22 min. 9.014 n 34 ” 8.33f " C -* 7.7711 ” 43 " 7.0524 « 25 « 6.47

30 " 27 * 6.2952 D 51 « 6.32

= . 144 * 48 * 6.29

■ :•'

D;

x 10”2

l i "2.4952.3522.202

35

2ABLE XIV

Solvent: Bioxanea: 0*10X955!. allyl alcohol For k: A: 43.945 B:For kl5 A: 0.1223 B: 27.9? 1r l B 6lT 5»

t V k x 10-39 hr. 38 min. 10.82 3.49021 *** 43... * 10*7544 ” 41 » 10.02 2.60071 " 30 n 9.90105 " 21 «

145 ■ 58 n9.289.12

1.462214 M 57 w 8.50 0.913312 « 13 ” 7.84430 * 25 w 6.98 0.277699 * 5 ” 5.821294 n 23 * 4.60 {eqnilihritto'

7.1613.2202.1261.6911.664

2. a: *0507481. allyl alcohol For k: A: 2447 B; .00094For kx: A: .09571 B; 17.91

h: .0498H. iodine Cs .003767 B; .003838 Cf 7.50 B: 9.26424

. t ' ’ V—

7 hr. 31 min. 11.7070 ” 2 ” 11.09191 " 35 w 10.01430 " 17 " 9.07838 • 54 ” 7.891078 ” 53 » 7.67At equilibrium 7.50

k z IQ-3

ar,1.3770.3410* M B0.219 2.327

Solvent: Propylene ohlorlfie1* s: 0.10441T. allyl alcohol

For ki A: 42.28 B: .0544 For ki: A: .099605 Bs 23.25

U.Iodine D: .007863 B: 9.44070

t V

12 hr. 36 min. 7.8530 " 38 ” 5.6347 ” 51 » 4.8960 ” 9 ” 4.5596 n 1 n 3.89193 " 37 " 3.18Final *.S0

kl = 10-3 _ k x io-3

3. as .00061!. Allyl alcohol For k « .0S03-.0037S7For k^: A: 0*03316^ 3: 5.181

12t

hr. 26 min.

v

11.3730 ” 50 n 9.6548 n. 16 w 8.7772 * 9 # 7.61

107 « 22 6.66217 " 24 * 5.14*09 " 12 * 3.58816 * 49 # 3.62

1: .04991% Iodine

C: 3.60 3; 9.71823

kX » 10-3

1.928as1.5821.4651.288

k % io-3

3.3113.1032.434

EEss

s » ff

2 Is

EggK

TABIS 371

1. a: 0.1008M. ollyl alcohol For k: A: 45,19 B: ,0509For kx: A: .08173 B: 19.53

Solvent: Isopropyl ether

t V

lir. 37 *la. 11*58" 56 " 10.09" 35 » 8.37n 41 « 4.97t. X3 » 3.56" 47 » 2.57

h: .049SM. IodineG: .003767 C: 2.36

B: .007609 D: 9.62394

k % 10-3 kl x lO"41.821 7.9881.335 5.9250.9404 4.577h i ! 7*5551.025 8.729

.049971’. allyl alcohol h: .05004%. Iodine

•000r kx: A: ,03808 B: 5*891 C: 4.22 B; 9.67861

t V k x 10~® kx x Uhr. 4 min. 12.08 3.164 14.95* 16 * " SO ”

10.689.40

1.7201.438 l : ? in 6 «. 7.40 1.233 6.070» 4 " 5.50 1.235 7.156* 18 " 5.09 1.245 7.894

TiBLS XVII

Solvent: !Totrachlorethane

1. a: 0.1052%. allyl alcohol b: .0459%. iodineFor k: At 41.59 E: .0553 C; .003767 D: .007540For kxt A: .09154 B: 21.65 Ct 2.65 Ej 5.48254

t v ' kx x 10**2 k x 10"2

hr,« 7 min. 6.03 1.762 . 3.571n 20 # 6.31 1.683 3.195* 34 # 4.61 1.486 2.728* 0 * 3.90 utat 2.04613 Tf 3.31 0,757 1.178n 56 # 2.76 0,496 0.4898it 0 # 4.21 .# 24 # 4.24

.05027%. allyl alcohol b: .05005%. iodine

■ * » •■"rdooieel— 1• kit A: .03033 B: 4.65 C: 3.40 2: 5.74126

t V ^ x 10-3 k % IP'Shr# 51 tain. 11.09 4.198 9.757

H 3 ft 10.18 3.634 8.521# 27 * 9.02 2.749 7.834# 47 # 8.11 1.968 4.375# 14 * 6,79 1.212 2.618« 61 * 3.36* 8 tf 3.40

3*

DisouBBion of Results

Be suits show that the reaction "between iodine and ally! alcohol in liquid solutions is catalyzed by the presence of little-ionised salts. The degfee of catalysis is nearly proportional to the amount of catalyst. Mercuric chloride was more effective as a catalyst than mercuric lotlie.Cadmium chloride was not as effective as mercuric chloride but was more so than mercuric iodide.

ftrom Tables III and VI it appears that the catalyzed reaction teem not give satisfactory balances for "k" or The reason for this is not apparent. It may be that the catalyzed and uncatalyzed reactions are of different order.Again other reactions may be proceeding, either as side chain reactions or parallel reactions, other than the one under in- veatigatien. This conclusion might be substantiated by the presence of a brick-red precipitate at times believed to be mercuric iodide. In this case again the equations used would not describe the reaction. In any event the mechanism of this reaction catalyzed is difficult of determination, if not impossible. Catalyzed reactions in general are among the unsolved problems of kinetics*8*

From the data in Tables VII to XVII inclusive it can be seen that the reaction goes to, or nearly to, completion in some solvents, while in others an equilibrium is reached. The

jU S B O S

40

most complete reaction wae otserve6 In the solvent carbon tetrachloride, liven in thlc Instance there t*as c very slight evidence of attainment of an equilibrium, ibis is farther shown by the calculated second order velocity constants* Three solvents gave satisfactory constants for "k", four gave satisfactory constants for nk i % two were nearly constant for nktr, and two gave neither a constant for *k" nor for These results are shown in tabular fora inTable m i l .

TABLE m i l

✓

Solvents riving satisfactory valuer f„k" "ki-

Carbon tetrachloride Absolute ethyl alcoholChloroform Absolute butyl alcoholCarbon bisulfide Propylene chlorideBensene , -Toluene

neither ”kn nor

leoprepylether

From tho grouping in Table m i l It will be seen that the one giving constants ,rkv for the reaction are those solvent* that have molecules with no polarity, or nearly so. Strictly, carbon tetrachloride is the only solvent that fits this requirement, for carbon bisulfide and benzene possess molecules that are unsaturated, and chloroform and toluene possess small electrical moments. It will be noted that the results in Tables m i , IX, X, XI, and XII indicate some negative catalytic effect for carbon bisulfide and benzene, and slightly greater negative effect for chloroform and toluene as compared with carbon tetrachloride.

Ibis 1b assuming the reaction in carbon tetrachloride to be standard, for, quoting Hoeluyn-Hughco,58 Mln the case of reactions uhleh cannot be measured in the gaseous phase, the rati® in such solvents cay be regarded as standard values with which rates in other solvents oay be ccopared.* While the addition of iodine to allyl alcohol may be mcaoureable in the gaseous phase, the investigation reported herein was solely in the liquid phciee. therefore it will bo aesuwsd legitimate

fZQto make such comparisons with the reaction under discussion.She research of some Investigators, particularly that of

W i l l i a m s , 41,%2 has shown that carbon bisulfide and benzene have very snail electrical moments, L’cAlpinc and Smyth-3 eoaslwded benzene had no electrical moment in the vapor phase. This is also true of chloroform^ and toluene,̂ 3 whose electrical moments are not great. A glance at Tables IX, X, XI, and XII will show little drift in °kn for these solvents, ana their values are not much smaller than for carbon tetraehloride. Although these solvents seem to have some negative catalytic effect, it in comparatively small. Therefore they may be considered as giving constants nkn, A comparison of "k" and for these solvents shows them tovary about the sane. Since those solvents In the first group of Table XVIII do give constants Mkn, it can reasonably bo concluded that the addition of iodine to allyl alcohol in these solvents is a second order reaction proceeding practically to completion according to the following equation:

4#

CH£:CH.CH2.0II -hiz— *> ch2i .chi.ch2.oh

She electrical ooneats of polar molecules rary with the phaccE in which they exist35»40 and with the solvent if they arc in solution. Ho references could be found which save the electrical moments of the molecules of the solvents herein considered under the same conditions. Therefore no true comparison of these solvents could be made in this respect.Ac stated in the Introduction, however, no general success has been won from generalising upon physical constants of the solute if or solvents.2 Such generalisations may hold for specific reactions but do not hold universally.

In the second column in Table XVIII are four solvents, aleolute ethyl alcohol, absolute butyl alcohol, propylene chloride, and tetrachlorothane, which give constants forthe opposed reaction. The constants for tetraohlorethane show some drift, but it le placed in this group. Undoubtedly other reactions Interfere in this solvent, for a pungent gas, believed to be hydrogen chloride, was detected occasionally. Such unknown complicating reactions make it difficult to determine the exact mechanism. Definitely, however, the reaction in these four solvents is reversible, and of second order in accord with the equation:

CHg:CH•CKg•OH+ Ig**^CHgl.CHI.CKg.OHAn inspection of the solvents in this group reveal them

to have molecules which possess marked electrical moments. This fact clearly differentiates thee from the solvents

43

listed in the first ccluan Trhieh give constants °kn for a completed reaction. A comparison of the velocity constants In fables YII, XIII, ZY, and XVII shows that butyl alcohol bes a positive catalytic effect while ethyl alcohol and tetra- ehlorethar.e seem to have a slight negative effect, fbe molecule of butyl alcohol possesses a higher electrical moment then does ethyl alcohol*44 In general, Smyth44 points out, the longer the carbon chain the greater the electrical mo- - meat* In the case of the reaction herein studied, the greater the electrical moment of the solvent molecule of chain compounds the more positive the catalytic effect. It must be kept in mind that this conclusion may not hold for other reactions eve* of a similar type.16

In the third column of Table XVIII are two solvents, dioxane and isopropyl ether, which So not give constants "k" or An inspection of fables XIV and XVI shows thattheir velocity constants drift downward. It will be noted that the molecules of these solvents also have polarity.The mechanism of the reaction in these solvents in difficult to determine, but it seems probable that the reaction Is reversible end second order but complicated by other reactions.

However, it may be that the nature of the reaction in these solvents has changed and cannot be described by the mathematical equations used or possibly not at all. The fact that dlorane46 and isopropyl ether4& can form oxoniua salts with iodine end possibly react with allyl alcohol may appreciably interfere with the reaction and cause the drift of the

44

constant.fhe reaction in fiioxane nan eloper than in any of the

other solvents trleti#The combined results of the reaction in the various

solvents chon that the reaction goes to completion in those solvents which have no polarity vihile an equilibrium is reached in solvents which possess polarity. The greater the polarity of the solvent molecules the greater the tendency of the reverse reaction. There is come tendency toward equilibrium* even in solvents which have no polarity. This is undoubtedly due mostly to the formation of complexes.47 Certainly the fact that allyl alcohol pcseeeses molecules with an appreciable polar!ty48 mould substantiate this conclusion.

no attempt will be made here to give a complete solution of the kinetics of this reaction. Attempts at such solutions in general have failed* particularly with respect to physical properties. 5*om the evidence obtained in these investigations polarity scorns to have an important bearing on the question# Just what hearing and to what degree will require much more investigation.

45

• Summary

The reaction of iodine with allyl aleoho1 in various solvents proceeds at a measurable rate.

Thie reaction can definitely be catalyzed by little ionized salts emeh as mercuric chloride, mercuric iodide, and cadmium chloride. It is not, however, autoeatalytic.

The mechanism of catalysis is not known.With carbarn tetrachloride, carbon bisulfide and ben

zene as solvents the addition of iodine to allyl alcohol gees to completion and is of the second order. The reaction in chloroform and toluene goes practically to completion, giving a fairly good constant Mk,t. With absolute ethyl alcohol, butyl alcohol, propylene chloride, and tetrachlorethane as thesolvent® this reaction reaches a definite equilibrium but

'

gives a second order reaction velocity constant for opposed reactions. The solvents dioxane and isopropyl ether did not give either constant *k* or nk^n«

In general, the reaction goes to completion and gives a second order reaction velocity constant in those solvents whose molecules have no polarity, and reaches an equilibrium and gives a second order reaction velocity constant for opposed reactions in those solvents whose molecules have

polarity.

BIBLIOGRAPHY

1. Lewie, Trans. Cheo. Soc.» 113, (1918), p. 471.2m Koolwyn-Rughec: Kinetics of Reactions In Solutions, p. 52.

y/ 3. Uilhcl^Y, Page, Annslcn, 81, (1850), p. 413.4. Woelwyn-H»gbGB: Kinetics of Reactions In Solutions, p. 85. 5* Boelwyn-Hu^ies: 1514., p. 52.6. I}* 5. Kharasch, I'.C, HcITat and P. R. ITayo: J. A. C. S.,

65 (1933), p. 2531.7. Coheni Organic Chemistry, 5th Edition, p. 1268. MoeIwyn^Eu^ies: Kinotice of Reactions in Solution, p. 68.9. C. H, Hinehelwood: Kinsties of Cheeioal Change in Gaseous

Systems, 2nd eS., p. 2.10. Perrin: Ann. fie Phyc., 11. (1915), p. 1

x/11. Lewis: Philosophy Haganine, 39, (1920), p. 2612. Lindesan: Trans. Faraday Society, 17, (1922), p. 658.13. Hinehelwood and Hutchinson: Proc. Roy. Soc., A, 111,

(1926), p. 245.. Hlnefaclwood and Thompson: Ibid., A, 113 (1526), p. 221.

Hlnehelwood and Aakey: Ibid., A, 115 (1927), p. 216. Hinehelwood: Ibid., A, 114, (1927), p. 84.

14. Hlnehelnood and Rusgrave: Proc. Roy. Soc., A, 135 (1932)p. 235.

15. Uoelwyn-Hughes: Kinetics of Reactions in Solution, p. 20.16. Moelwyn-Hughes: Ibid., p. 2.

47

17, XfoeIwyo-Hugbea: Ibid., p. 52, footnote.18. Bray ond llvingeton: J. Acer. Chem. Soa., 45, (1523), p. 1654.

Bray anfl Llvlngobon: Ibid., 50,(1928), p. 1654.Livln^eton: Ibid., 48, (1926), p. 53,

19* Boanon and Le Roecignal: (Trane. Qica. See., 85, (ISOS), p. 70S.

20. Borrieh: C. A. 18, 493. J. C. S. 123 (1923), p. 3006.21. Heieigs C, A. 27, 1860; J. A. C. S. 55 (1933), p. 1297.22. ritettkuri; C. a . 28, p* SSf,23. Uilliamo: C. A* 26, p. 4525; J.A.C.S. (1932), p. 1747.24. Hoafenn and Kiarsuthers Berichto 42 (1910), p, 4483.25. Stewart end Edland2 C.A. 18, p. 493; J.C.S. 123, (1923)

p. 3006.26. WilUcao and Jamec: A. 22, p. 1968; J. C. 2. (1528),

p. 343.27. Roberteon. Glare, Bcbaught and Paul; J. C. S. (1557), p. 355* 26* niausG and Sdu.11; C.A. 22, p* 221,

48

38• MoeIwya-Hughcc: Kinetics of Reaction# in Solutions, p* 71.39. Vllliaao and Krclima: J.A.C.S. 49 {1927},pp. 1676, 2408.40. VlUIaBn and Ogg: Ibid., 60 (1928), p. 94.41. bllllnmE and Sohelngel: Ibid., 50 (1928), p. 362.42. X/illiamo and X/elecberger: Ibid., 50 (1928)e p. 2332.43. Kc A Ip In® find Stoyth» J. A. G. S. , 55 (19 So 5, p. 453.44. Smyth: J. A. C. S. 46 (1924), p. 2151.45. Whitmore: Organic Cheolctry, p. 875.46. Whitmore: Organic Chemistry,pp. 152-3.47* 1;oeIwyn-Hnghon: Kinetiefi of Reactions in Solutions, p. 148.

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