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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.
Rights Copyright © is held by the author. Digital access to this
materialis made possible by the University Libraries, University of
Arizona.Further transmission, reproduction or presentation (such
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Download date 05/04/2021 14:31:24
Link to Item http://hdl.handle.net/10150/553407
http://hdl.handle.net/10150/553407
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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-'
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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
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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
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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
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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
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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.
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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.
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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.
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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
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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
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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*
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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
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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
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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
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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
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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
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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.
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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
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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.
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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
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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
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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.
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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
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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
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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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