«V' '« 'VIA
:.:X
i IIM iY.'hn, HO/1
O-i
1(
5'il"nir,.
'»!
Txffi DIFFERENCE IH CHEUICAL ACTIVir/ OP
FREE AOT) SEMIC0I.3INED ^ATER
AS ILLUSrRATED 3Y THE EFFECT OF
NEUTRAL SALTS ON THE HYDROLYSIS OF ACETIC ANHYDRIDE
DISSERTATION
Submitted to the Board of University Studies of
The Johns Hopkins University
in Conformity with the Requirements
for the Degree of Doctor of Philosophy
by
GERALD CHARLES CONNOLLY
June 1916
-i- i «-f
ACKN07LED(}M1-2IT.
The author accepts thia opportunity to express his ;^rati-
tude to Professors luOrse ani Rem8en;to Associate Profeesors
?razer,Reld and Lovelace ;to Collegiate Professor Gllpln;and
to former Associate Professor Acree.for advice and instruc-
tion which have been at all tines cheerfully given both in
the lecture room and in the laboratory.
He also feels under obllflrations to Dr. Davis and Dr. Lloyd
for many valuable suggestions and criticisms.
Especially is the writer Indebted to the late Professor
Harry C. Jones, under whose supervision and untiring care
this investigation was carried out, and whose untimely death
means the loss of both a teacher and a friend.
TABLE OP con TENTS.
Page
Aoknowled.^ent ^
Introduotlon ^
HydrolyslB ^
Hydrolysis of Acetic Anhydride .... 5
Hydrolysis of Salts H
Neutral Salt Action 13
Statement of Problen 18
"Tperlmsntal 20
ruriflcation of Acetic Anhydride . . 20
Purification of Salts 22
Apparatus 22
Solutions 23
liethod of Procedure 24
Calculations 26
Data 26
Tables ■ 27
Discussion ^'^
Biography ^^
THE DIFFERENCE IN CHEMICAL ACTIVITY OP
FREE AND SEMICOMBINED TJkTER AS ILLUSTRATED 3Y THE
EFFECT OF NEUTRAL SALTS ON THE HYDROLYSIS OP ACKTIC ANHYDRIDE.
The present Investigation was undertaken for the purpose
of obtaining further evidence on the solvate theory of solu-
tion. The solvate theory was proposed in this laboratory by
Professor Harry C.Jones in 1900, to account for certain ap-
parently abnormal freezing-point lowerings produced by cer-
tain electrolytes when dissolved in water. Prom this begin-
ning the work has expanded in many directions until there are
at present sixteen distinct lines of evidence, all pointing
to the validity of the view that when a salt is dissolved in
a solvent there is more or less combination between the salt,
or the ions resulting from it, and the solvent. The magnitude
of this solvation depends upon the nature of both the solvent
and the dissolved substance .
HYDROLYSIS.
The term hydrolysis is applied to a number of chemical re-
actions in which there is first the addition of water to a
complex, and then the decomposition of the product into sim-
pler substances. Prom this definition it is evident that the
reactions included under hydrolysis are numerous and varied.
There are In general four main dlvlalono of hydrolyela:
(1). HydrolyslB -if metallic salts.
(2). Hydrolysis of esters and closely associated subntancee,
such as amides, nitrlles.aold chlorides, acid annydrldea ,eto.
(3). Hydrolysis of complex carbohydrates and ^luoosldes.
(4). Hydrolysis of polypeptides and proteins.
In this discussion we will confine ourselves almost entirely
to the first two divisions, for these are the only forma of hy-
drolysis which come within the scope of this Investlf^ation.
Hydrolysis of Acetic Anhydride.
The hydrolysis of acetic amhydride has been studied by sev-
eral Investij^ators with varying degrees of success. The term
"hydrolysis of acetic anhydride" is used here in preference to
the term ''hydration of acetic anhydride" used by other investi-
gators, since it is more in accordance with the definition of
hydrolysis previously stated.
Menschutkin and VaEilieff , in studyin.^ the decomposition
of acetic anhydride by water, attempted to find a constant for
the velocity of hydrolysis. They did not succeed because the
two substances were not mlscible in all proportions, and no sol-
vent could be found that did not act either on the substances
themselves or on the products of the reaction. They studied the
change In a homogeneous mixture of acetic acid, acetic anhydride
and water, and found that the reaction was not as rapid as had
generally been supposed. With a mixture of equal portions of
1. Journ. Rusa.' Phyp. Chem. 3oc. ,21,188 (1889),
acetlo anhydride and water at 19° only about on«-half the anhy-
dride wae hydrolyzed at the and of bIx hourn, and eleven dayn
were neceeaary for complete hydrolysln. In the followln/r table
a oompariaon la made betwoen the velooltlea of deoompoaltlon of
acetio anhydride, acetamlde and ethyl acetate by one ;^am-equlv-
alent of water at 100° under the same conditions. Phe experi-
ments were carried out In the presence of acetlo acli.
Substance
Acet .Anhyd.
+ 1 H2O
Acetamlde
+ 1 HoO
Kthyl Acetate,
+ 1 H2O
Acetic ..cii
Added
11.86 ^
15.85 '^0
11.45 "
Time
Decompoaition In percent.
1 min.
11 "
61 "
121 "
181 "
1
25.68
83.9
98.5
99.5
99.7
4.51
4.64
4.94
5.82
6.41
0.2
0.5
0.87
0.99
\-
The acetic anhydride was almost entirely deoompoael at the
end of one hour, while the decomposition of the acetamlde was
slight and that of the ethyl acetate had hardly beprxm.
Hinaberg showed that acetic anhydride is soluble in water
in the anhydride state, ajid that the ijresence of water la not an
ohstacle to the employment of the anhydride as such.
A. and L. Luralere and 3arbler found that a solution of
2. Ber.d.deutach.chem.GeBell..23,2962 (1890).
3. Bull. 300. Chlm. (Ill) 33.783 (1905).
aoetlo anhydride In water poBBeasee praotloally all the prop-
erties of acetlo anhydride and 1b oufflolently atable for
acetylatlon purposes. 3y simply shakln*? the anhydride with
water, they obtained a twelve per cent, solution, whloh read-
ily aoetylated amino- and Irnlno-oompounds.
4
In a second paper A. and L. Lumlere and Barbler showed
that if more than twelve parts of the anhydride are used, solu-
tion Is Incomplete, They prepared five euid ten per cent, solu-
tions of acetlo anhydride In cold water and set them aside.
From these solutions aliquot parts were withdrawn every ten
minutes and added to a alight excess of aniline, the excess be-
ing known. Reaction took place quantitatively between the ani-
line and the acetlo anhydride not hydrolyzed by the water, with
the formation of acetanlllde and an equivalent of acetlo acid.
The total acid present was then determined by titration with a
normal solution of sodium hydroxide in the presence of phenol-
phthaleln. Prom this the degree of hydrolysis of the acetlo
anhydride was calculated, and the results obtained are given In
the following table. (See page 8).
Prom these results it can be seen that the rate of hydroly-
sis is fairly rapid at first and then gradually decreases. It
Is the more rapid the greater the initial dilution of the anhy-
dride and the higher the temperature.
Alcoholic solutions of the anhydride were also prerarei,
and it was found that when molecular proportions were used, es-
terlfloatlon was Incomplete, even after a month.
4, Bull. 300. Chim. (Ill) 35,625 (1906),
HYDROLYSIS OP ACETIC ANHYDRIDE BY '.VATER.
Time
10
20
30
40
60
60
70
80
90
100
110
120
140
16C
5 "' Solution
1.^ T
9.2
52.5
74.2
89.7
95.7
100.0
GO
4.6
35.0
48.4
60.8
69.0
76.2
8C.4
85.5
89.6
93.8
96.9
100.0
10 % Solution
1 r, ' -
11.5
58.2
71.0
78.9
86.6
91.7
93.3
94.6
96.4
97.9
100.0
9.8
34.6
51.1
60.0
67.0
73.3
77.9
81.5
85.1
88.9
92.8
94.8
95.8
98.5
lOC.C
8
5
Benrath , by meame of change in density, attempted to
measure the rate at which acetic anhydride combines with water,
using acetic acid as the solvent. He concluded that, with
equivalent ouantltles of anhydride and water, the reaction was
monomolecular. I.e., It depended on the concentration of only
one of the reacting substances.
Rivett and Sldgwick , using dilute aqueous solutions, fol-
lowed the hydrolysis by determining the electrical conductivity
at 25°. The velocity of hydrolysis showed that the reaction
was monomolecular and was not catalyzed by hydrogen ions; and
that beyond a certain point the constant decreased steadily
with Increasing concentration. From this they concluded that
the mechanism of the hydrolysis of acetic anhydride is quite
different from that of the hydrolysis of esters and amides, in
which case the velocity constant is proportional to the concen-
tration of the hydrogen ions.
7
Orton and M.Jones ,on the other hand, consider the hydro-
lysis of acetic anhydride to be entirely analogous to the hy-
drolysis of esters and amides. They conclude that the hydroly-
sis of acetic anhydride in acetic acid as the solvent is a slow
reaction of the second order; that on dilution the increase in
velocity is approximately proportional to the amount of water
present, and that the relation of the velocity factor to the
temperature is normal. The effect of catalysts was also studied.
5. Zelt.physlk.Chera. 67,501 (1909).
6. Joum.Chera.Soo. 9^.733,1677 (1910).
7. Journ.Chem.Soc. 101,1706 (1912).
10
It wns found that aclde are powerful catalyntn of the hydroly-
al9. The effect 1b moat noticeable In media oontalnln>r but
little water, and dlrainlshea as the proportion of the water
Increases, belnf? least obvious In pure water. The value of
the velocity factor Is a linear function of the concentration
of the acid. Alkalies and hydrolyzed salts were also found to
act as strong catalysts of the hydrolysis In aqueous solutions.
The following equations were trlven to represent the raeohanlam
of the hydrolysis:
(I) AcgO + HgO = 2AcOH
(TI) Aco0-«-H20+r s aAoOH + H"^
(III) AOgO+HgO+HX = 2AO0H+HX
(IV) ACgO-t-HgO +0H' = 2AcOH +-0H'
Any one of the four forms could predominate according to
the conditions, medium, etc. In aqueous solutions the choice
lies between (I), (II) and (IV).
Philip made a study of the reaction between acetic anhy-
dride and water in glacial acetic acid by determining the
freezing-points of the mixture at freouent intervals.
9
7ilsdon and Sidgwick studiei the rate of hydrolysis of a
numbers of acid auihydrldes. The rate of formation of the acid
was determined by measuring the increase in the electrical
oonduotivlty of a solution of the anhydride In water.
Verkade also made use of the oonduotivlty method. Hie
8. Proc.Chem.3oc., 28,259.
9. Journ.Chem.Soo. 103,1959 (1913).
10. Rec.Trav.Chlm. 35,79 (1915).
11
results for the hydrolysis of aoetlo anhydride a^jree very
closely with those obtained by Rlvett and Sldtrwlok ^.
Hydrolysis of Salts.
It is a well-knovm fact that oertnln salts, even thoup^h
they contain the strictly equivalent quantities of acid and
base recuired for 'neutrality"', when dissolved in water are
not neutral to indicators but react either acid or alkaline,
This was first noticed by H. Hose in workinicj with certain
basic salts, but was not explained satisfactorily until
Arrhenlus proposed his theory of electrolytic dissociation.
In the light of this theory acidity is due to the presence of
an excess of hydrogen ions, and alkalinity to the presence of
an excess of hydroxyl ions. These ions cannot be accountei
for by the salts themselves, therefore they must be accounted
for by the water.
7ater must contain both hydrogen and hydroxyl Ions. rhe
Ionization constant of water can be calculated by the equatlon-
H^X OH ', y.
H2O
Since the active mass of the nonlonlzed water is so great in
comparison with the active mass of the ions, it may be consid-
ered constant, '.ye then have -
H*x 0H-= k^go
the value of k being 1.2 x 10-14 at 25°. rhls ionization is
the same in all aqueous solutions. The value kj. ^, however,
Increases with rise in temperature. This Increase is most
11. Pogg.Ann. 83.132,417 (1861).
It
probably due to the breaking down of the anaoclated moleoulea
intc the simpler ones whloh are more easily dlseoolated. Pure
water contains an equal number of hydropfen and hydroxyl Ions,
and therefore must react neutral. Ii'urthernore, thin relation
holds for any neutral solution. To be acidic, a solution must
contain an excess of hydrof^en Ions; to be basic, an excess of
hydroxyl Ions. To determine, then, whether a solution is neu-
tral or not we make use of indicators, such as litmus, methyl
oran(?e, phenolphthaleln, which give evidence by their color
changes.
When a normal salt is dissolved in water partial hydrolysis
takes place, yielding free acid and free bane. V.'hether the
solution will react acid or alkaline will depend on the le/rree
of dissociation of these products of hydrolysis. If follows,
therefore, that there are four types of salts which may under-
go hydrolysis: (1) salts derived from strong acids and strong
bases; (2) salts of weak acids and strong bases; (3) salts of
strong acids and weak bases; (4) salts of weak acids and weak
bases.
All salts except those of the first type are hydrolyzed to
a considerable extent, due to the small degree of dissociation
of one or of both of the products of hydrolysis. Salts of
strong acids and strong bases under ordinary oonditione do not
undergo hydrolysis.
The determination of the degree of hydrolysis is not aooora-
plished without difficulty. The free acid or base cannot be
directly titrated with a standard solution, for equilibrium
would bo destroyed at once and neutrality would be reached
13
only when the salt was oorapletely cleooraponed, A method nmot
then be employe I whloh will not destroy the hydrolytlo eqall-
Ibrlura. The method? most generally usel are: (1) the deter-
mination of the velocity constant for the hydrolysis of an es-
IJi
ter , for this Is proportional to the amount of free acid or
alkali present; {2] the determination of the rate of Inversion
12
of cane sugar ; (3) the determination of the electrical con-
ductivity of the solution ^^; (4) the determination of the
12
coefficient of distribution between two solvents . In addi-
tion to these there are also many other methods of more or
less limited applicability.
Only those salts were used in this investigation whloh
were proved by the above methods to be nonhydrolyzel. ?or
these salts the values of k calculated according to the equa-
tion -
(salt) _ k
(Acid) X (Base) ^^HpO
are so small that they need not be taken into account. The
salts were further tested according to an observation made by
13
Salm , that salts which give no reaction with litmus have a
concentration of :i^and OHlons less than 1 x 10'^, a value so
small that It Is negligible.
Neutral "alt Action.
In a discussion of neutral s«\lt action one must distin-
guish clearly between the effect produced by a neutral salt on
12. R.C.Parmer, 3. A. Reports, 240 (1901).
13. "^eit.physlk.Chem. 57,471 (1907).
14
the catalytic activity of an aoid (or alkali), and tha affaot
of tha neutral salt on hydrolyols by water alone. It 1b the
latter effect In which we are moat Interested In this Investi-
gation, although the former 1b what la (generally unlerBtood by
the term "neutral salt action".
Effect of Neutral Salts on the Catalytic Activity of Acids.
14
It was early found that the addition of a substance which is
lar<?ely ionized in aqueous solution alters the rate of hydro-
lysis of esters or of carbohydrates by stronf? acldp. Hhitt has
been proved by the addition of metallic chlorides to mixtures
in which hydrochloric acid is the catalyst, the addition of
bromides to hydrobromic acid, and of nitrates to nitric acid.
Those chlorides which are highly dissociate 1 have much the
same effect, while a salt like mercuric chloride which Is only
partially ionizel has a much feebler action. Non-electrolytes
such as the alcohols or sus^ars have but little effect on the
hydrolytlc activity of the hydrofren ionn.
The action of the neutral salt is not always to accelerate
the hydrolysis; often there is a retardation. There are also
well-defined differences between the Influence of neutral salts
on the rate of inversion of cane sugar in the presence of acids
and their influence In the catalytic hydrolysis of esters.
The velocity of the inversion of cane aup'ar Ip increased to a
much greater extent by the addition of certain concentrations
of salts than is the velocity of the hydrolysis of esters.
Neutral salts have in general a retarHnr- effect upon the
14. Journ.prakt.Chem. 86,381,401 (1862).
16
hydrolysis of •et^ro and amld«n by alkallnn. S«nt*r ■^^, how-
aver, found that th« hydrolysis of sodium ohloroac«tat« by
sodium hydroxld* was graatly aooel«rat«d by tht presence of
ntutral salts. It has b«en shown that neutral salt action Is
independent of the oonoentratlon of the compound hydrolyted;
Is prOTOrtlonally ji^reater the more dilute the ncld solution;
Is not (greatly Influenced by teraperatur« or pressure; and is
Independent of the nature of the sold employed as catalyst.
In addition, Poma has deternlned that the intensity of
the action developed by neutral salts bears a strict relation
to the chemical nature of the ions of the salts, and dirainiahes
in passing from chlorides to bromides to nitrates to iodides,
in succession; that It is independent of the chemical nature of
the oations; and, finally, that It seems to be proportional
not to the concentration of the salt in the solution, but to
the concentration of the ions.
Effect of Neutral 3alt3 on Hydrolysis by Water Alone.
Probably the first work done on neutral salt action in the ab-
17
sence of an acid was by 3nlth , who investigated the effect
of neutral salts on the rate of inversion of cane sugar. He
found that salts of weak acids had almost no effect, while po-
tassium chloride and sodium sulphate, the more nearly neutral
salts, had considerable effect.
Senter showed that neutral salts have practically no effect
15. Journ.Chem.oOc. 91 ,473 (19C7).
16. Medd.K.VetenskapBakad.Nobellnst. ,2, No. 11. pp. 1-28,
17. -^elt.-nhvHlk.'^hem. 25,144 (1898).
16
on the deoomposition of aodlum ohloroacetate by wftter.
18
K«llo;5(^ studlei the effect of the neutral naltB, potaa-
slura ohlorlde, potassium bromide and potassium lodlie on the
velocity of .the hydrolysis of ethyl acetate. The reactions
were carried out in sealed tubes at 100°, usin;^ a fixed quan-
tity of ester and varying: concentrations of the salt solution.
The results obtained show that the specific influence of salts
is ^Treater in somewhat dilute solutions. As the concentration
is increasei, the effect gradually becomes less; until it
reaches zero, and then becomes nef^atlve in character; for exam-
ple, a 4-norTnal solution of potassium chloride hydrolyzes the
ester more slowly than pure water itself, Kellogg found a de-
crease in the accelerating power from chloride to bromide to
iodide, which is in reverse order to their stability,
19
Henderson and Kellogg continued the investigation using
the chlorides of sodium, lithium, calcium, strontium and bar-
ium, and the chloride and iodide of cadmium. They carried out
the work under the sarre conditions as before, and in addition
measured the conductivities and viscosities of the solutions
at the concentrations and temperatures employed in the experi-
ments, and from these calculated the degree of ionization.
They found that the salts which produce the greatest effect
are those which are the least ionized. The accelerating effect
of lithium chloride is greater than that of sodium chloride
although the degree of ionization of the former is less; while
18. Journ.Amer.Chem.Soc. ^,403.886 (1909),
19. Journ,Amer,Chem,Soc. 35,396 (1913).
17
the chlorides of oalolum, barium and ntrontlura hav« a grtatar
«ff«ot than either sodium ohlorlde or potaaelum chloride, al-
though they too are less lonlTied. Cadmium ohlorlde, the least
Ionised of all the chlorides studied, produced the (greatest
effect, due probably to the hydrolysis of the salt. Henderson
and Kellogg concluded that the effect produce 1 by n neutral
salt on the hydrolysis of ethyl acetate Is due to a specific
influence on the non-lonlzed portion of the salt, rather than
to any function of the ions.
Holmes and Jones ^^, and later Lyons, Connolly and Jones,
for reasons to be discussed later, took up the study of the
influence of neutral salts on the rate of hydrolysis of methyl
acetate and methyl formate by water. They used three concen-
trations of salt solutions: molar, half-molar and quarter-molar,
and found that with a few exceptions all the salts studied
have a pronounced accelerating effect upon the hydrolysis.
They also found that there is a concentration of the neutral
salt at which the effect is a maximum.
There have been several sup^rrestions put forward to explain
21
neutral salt action. Arrhenius " proposed that the salts nay
affect the substance which is being hydrolyzed; that there nay
be present in the solution an equilibrium between an active
and an Inactive form of the substrata, and that this equili-
brium nay be altered through changes of temperature or ionic
concentration. Armstrong and Caldwell concluded that the salts
20. Carnegie Inst. ^ash. Pub. No. 230 (1915).
21. Zeit.physik.Chem. 4,226 (1889).
18
act by re-.oving part of th« water In th« form of definite hy-
drated oompounde, and In this manner Increnpe the oonoentratlon
of the reaotinflr subatanoe. Stle^^llts? explained salt effect In
(general by the theory that the presence of anlts In the eolu-
tlon Increases the dielectric constant, or, at any rate, the
ionizing power of the solvent. All of these t/ieoriea are
plausible, but it la highly improbable that neutral salt action
is due to any one cause exclusively.
STATEMENT OF THE PROBLEM.
Anderson and Jones ^^ found that the absorption spectra of
solutions of salts like neodymium chloride and nitrate, in non-
absorbinj? solvents like water and alcohol .depend largely on
the nature of the solvent in which the salt was dlssolvedj for
example, neodymium chloride dissolved in water has a different
spectrum from that of neodymium chloride dissolved in methyl
or ethyl alcohol. They found for the first time what they
called 'solvent bands', showing that the dissolved substance
was combined with more or less of the solvent forming in the
one case "hydrates" and in the other "alooholates". Alcohol-
ates had, as would be expected, very different resonance from
hydrates.
Strong and Jones ^^ extended the work of Anderson and Jones
to a larp:e number of solvents and to a fairly large number of
non-absorbing salts, vid showed that these advents had a marked
23. Carnegie Inat .'Yash.Pub.No.llO (19C9).
23. Carnegie Inst.-7aBh.rub.I.'o.l30 (1910); IIo. 160 (1911).
19
influ«no« on the absorption Bp«ctra of th« ealte dlBBOlv«d In
them. They were able to dlatinfirulsh between the spectra of
salts dissolved In a normal alcohol and thone In the Isomeric
alcohol.
24
Guy and Jones built the most sensitive radlomlcrometer
constructed up to that time, and by means of it they studied
quantitatively the Intensities of absorption lines and bands.
They found that while solutions of slif^htly hydrated salts were
about equally transparent with pure water, solutions of stroni^-
ly hydrated salts were very much more transparent than pure
water.
The work of Guy and Jones was repeated by ."hneff er.laulus
and Jones , using an even more sensitive radiomicrometer
constructed by Shaeffer; and they confirmed the conclusions
reached by Guy and Jones. They found solutions of strongly
hydrated salts which were as much as 40 per cent, more trans-
parent than a depth of pure water equal to the water in the
solution in question. This showed that water of hydration has
a very different resonance from pure water.
Havinff found this physical difference between combinel and
free water, the problem wag then to see if there was any chem-
ical difference. Therefore, Holmes and Jones ^0 took up a
study of the action of strongly hydrated salts and slightly
hydrated salts on the hydrolysis of methyl acetate and methyl
formate. The method used consisted In measuring the velocity
24. Carnea-ie Inat .7nBh.Pub.No.l90 (1913).
25. Carnei^ie Inst, ./ash, Pub. No. 210 (1915).
20
of hydrolyalB of the enter by pure water and by solutlono of
slightly and otronRly hydratel aalte. The solutions were pre-
pared In suoh a way that the amount of water In each wan the
same, and was enual to the amount of pure water employed,
raking Into account the hydrolysis of the stron<?ly hydrate!
salts, they found that these snljjs hydrolyzei the ester much
more rapidly than pure water Itself.
The reaction studied by Holmes and Jones was a very alow
one and indioatei that combined water has f^eater activity
than free water. 'Ye wished to Investigate the same problem,
using a reaction that proceeded much more rapidly; therefore
we chose the reaction involvinf? the conversion of acetic anhy-
dride Into acetic acid.
EXPERIMENTAL.
Purification of Acetic Anhydride.
Pure acetic smhydride was necessary for the work. The
physical properties as described in the literature vary p:reit-
ly. The boiling-points given range anywhere from 135* - 140°,
at 760 mm. pressure. The densities given vary between 1.07
and 1.09. Prom this it can be seen that it was impossible to
test its purity by the ordinary simple means. Acetic acid is
the impurity most likely to be present in the einhydrlde, emd
is very difficult to detect if only small amounts are present.
0.51 pram of pure acetic anhydride, when completely hydrolyr-el,
is equivalent to 100 cc. n/10 solution of sodium hydroxide;
while the same weight of a mixture containing 1 per cent, of
81
aoetlo acid la oruivalent to 99.85 oo. this la within the ex-
perimental error.
Methods of finding the actual percentage of acetic aold
and amhydrlde In a mixture have been f^lven by Plckerlnp- ^^,
Mensohutkln and Vaellleff ^'^ , Preadwell ^^, Kdwarda and
Orton " , «md Orton and Jones ^ , Pickering determined the
freezing-points of the solutions of anhydride and water, and
compared them with the freezing-points of knovra concentrations
of acetic acid. luenschutkin and Vaallieff treat with aniline
and water, and determine the acidity after the reaction -
CgHgirHg +(01300)30 :: CgH5NHCOCH3 -f-CHgCOOH
has taken place. Treadwell recommends treatment with barium
hydroxide solution and titration of the excess of the latter;
while Jldwards and Orton convert the anhydride into acetanillde,
the latter into phenylacetylchloramine, and then determine the
chloramine volumetrically.
The method finally adopted to purify the acetic anhydride
was that of repeated dlBtillation, using a five-bulb distilling
head and discarding the first and last fractions. Phis gave
an anhydride which distillel practically constant at 138°-139°,
Specific f^ravity determinations, using a 10 cc. pyonometer,
gave a mean value of 1,0852 at 15°/4°. ?he acetic anhydride
26. Journ.Chen.:ioo. 63,1000 (1893).
27. Journ.:^88.Phys.Chem.Soo. 21,190 (ISBg).
28. Analytical Chemistry ,1914, Volume II.
29. Journ.Chem.Soc. 99,1181 (1911).
30. Journ.Chem.Soc. 101,1720 (1912).
22
was further tested by titrating weli^hed samples directly, and
by titrating weighed samples using the method advocated by
Mensohutkln and Vasllleff,
Purification of Salto.
Only the purest salts obtainable were used. They were
usually Kahlbaura preparations, although some of other well-
known firms were used.
These salts were dissolved In conductivity water, filtered
from any foreign matter present, and then recrystalllzel one or
more times.
Apparatus.
Thermostats. The constant-temperature baths were of
31
the Improved form designed by Davis of this laboratory.
These baths are substantially constructed of copper and are of
about 60 liters capacity. Gas is used for maintaining con-
stant temperature, the heat being applied directly to a heavy
galvanized iron pipe outside the bath, through which a part of
th<; water is kept circulating by means of the propellers.
Only a small portion of the water in the thermostat comes into
immediate contact with the heated surface, and this portion is
subsequently mixed with the main body of water, thereby secur-
ing much more even distribution of temperature. In these baths
the temperature can be kept constant well within C.08°,
The thermometers were of the differential 3eokmann type.
They were compared with a standard thermometer, which had been
31. Carnegie Inst .7ash. Pub. IJo. 210 (1914).
23
oallbratei at the Bureau of Standards.
Plaaka, pipettes and burettes for measuring purpoaes were
all carefully calibrated by weight.
All bottles used (varying In content from 50 oo. to 6 1.),
and all measuring flasks were of Jena glass.
A special apparatus vms used for the alkali solution, to
protect It from carbon dioxide and water vapor In the air.
Solutions.
The water used In the preparation of the solutions was
purlflei by the method of Jones and Mackay ^^ as modified by
33
Schmidt . It had a conductivity at no time fpreater than
2 X 10-6.
The aniline used to combine with the excess of acetic an-
hydride was the purest obtainable. It was further distilled as
many tines as necessary to remove all decomposition products.
The slightly colored product was then kept in a cupboard pro-
tected from light.
The solutions of the non-hydrated salts were made up dl-
thope
rectly by weight, while/of the hydrated salts were analyzed
gravlmetrlcally and diluted to the requlrel strengths. The
chlorides of barium, strontium, calcium and magnesium were de-
termined as silver chloride, and the sulphates of sodium and
magnesium were determlnel as barium sulphate.
The solution of sodium hydroxide used in titrating the
32. Amer.Chem.Journ. i7,83 (1895).
33. Amer.Chem.Journ. 19,90 (1097).
24
aoetio acid formrd oy tr.e nydrolyala of the acetio anhydride
was made up approximately half-normal, uslnf? "aodlum hydroxide
from alcohol". It was preserved in an apparatus proteotel
from the Impurities in the air. It was standardised by titra-
tion aprainst a««lnirt a solution of sulphuric acid of about the
same streni^th (0.4115 N.). The sulphuric acid had been stand-
ardized as barium sulphate.
The indicator used was phenolphthalein, as it f?ives the
best results in titratini^ a weak acid with a strong alkali; the
only objection bein/? that it is also sensitive to carbonic
acid. Corallin had been tried, but was not so satisfactory.
Method of Proceiure.
The method in principle is a modification of that of
Menschutkin and Vasllleff , and later employed by A. and L.
1 4
Lumlere and Barbler • . In order that the results should be
comparable, the amount of water present must be kept constant,
therefore the specific gravity of the salt solution was first
taken, ^ivln^ the welprht of 1 cc. j'rom analysis that part of
the weight due to the anhydrous salt alone was known for each
cubic centimeter. This known weight of salt, subtracted from
the weight of 1 cc. of solution, t^ave the weight due to the
pure water alone. This, divided into the weight of 1 cc. of
pure water at that temperature, gave the amount of solution In
cubic centimeters eouivalent to 1 cc. of pure water. Phe
amount of solution thus calculated was pipetted into a 250 oo.
Jena bottle. An equivalent of IOC oo. of pure water was taken
in all determinations. The bottle was suspended in the
26
oonatant temperature bath. There wan also plaoel In the bath
a bottle containing the anhydride and a number of small empty
bottles of 5C 00. capacity.
Whan all had oome to the temperature of the bath, the bot-
tle wag removed and 5 oo. of the anhydride Introduced. Time
was reckoned from when the anhydride was first addei. Solution
took place Immediately on shakln(^ except In the cage of the
▼ery concentrated solutlona. Eoual allruot portions were re-
moved and placed In the small 5C oc. bottles, the whole being
kept In the bath. These small bottles were removed, first ev-
ery 5, then every 10 minutes, and a slight known excess of
aniline added. On shaking, this combines with the residual
acetic anhydride, precipitating acetanlllde and liberating an
eculvalent of acetic acid. In one bottle of each series the
reaction was allowed to go to completion without the addition
of aniline, so as to control the results obtained.
The total amount of acetic acid was then determined In the
bottle by titration with the half-normal solution of sodium
hydroxide In the presence of phenolphthaleln as indicator.
Never less than 10 oo. nor more than 25 cc. of alkali, as
measured in a 50 cc. burette, were required to neutralise the
acetic acid.
Two temperatures, 15° and 25°, were employed. Only one
of
ooncentratlon/acetlo anhydride was used (approximately 5 per
cent,;, because if two were employed the results would not be
comparable on account of volume changes. For the salts molar,
half-molar and quarter-molar solutions were taken In all oases
and whenever possible solutions of greater ooncenrracion.
26
Mtaouretntnta of the velocity were not taken for longer
than 60 minutes at 15° and 40 mlnutea at 25°, for It wae found
that the hydrolynln of the acetic anhydride by water was then
praotloally ooraplete.
CALCULATIONS,
Pron the total amount of acetic acid, as determine* by
titration with the alkali, that due to the water alone must be
calculated. The simple formula -
y = 2z - X
Is used, where y la the amount of acetic acid due to the water
alone, z is the total amount of acetic acid measured by titra-
tion, and X Is the total amount of acid that can be formed If
all the acetic anhydride has been hydrolyzed.
The results obtained for the "control" bottles, when sub-
stituted in the formula, should give the same values for x and
y, which would be equivalent to 100 per cent, hydrolysis.
DATA.
In the following tables the concentrations of salt solu-
tions are !^, molar; I'/s, half -molar, etc. Time Is expressed in
minutes. All results are expressed in percentages, 100 per
cent, meaning complete hydrolysis of the acetic anhydride. In
each table there is placed for comparison a column showing the
percentage decomposition of acetic anhydride by water alone.
27
Table I.
Potassium
Chloride
at 15°.
I'/ater
3 M
Concent
2 M
ration
rim*
1 *
1» /o
"A
5
30.99
14.68
81.71
26.7;^
^v.c;
ox.tia
10
54.15
30.16
38.96
48.85
50.07
53.55
20
78.43
49.67
59.93
69.72
7S.51
77.33
30
9C.22
63.67
73.57
82.70
86.53
89.16
40
96.87
73.70
83.09
90.76
93.21
94.32
50
98.18
81.65
89.24
94.51
96.61
97.48
60
99.01
86.54
92.36
97.08
98,13
98.63
Table II.
Potassium Chloride at 25°.
Concentration
rime
V/ater
2 U
5
44.54
24.82
27.58
36.56
1 42.18
44.14
10
72.76
45.68
55.12 64.98
! 68.88
71.98
20
93.71
70.82
£14.63 , 87.87
91. C8
92.16
30
98.31
84.63
91.50 95.99
96.99
97.59
40
99.63
91.86
1 94.r'T 98.13
1 98.64
99. CO
26
Tlrae
Water
Pable III.
Sodlun Chloride at 15°,
Concentration
4 H
3 H
Z M
U
m/2
M/4
5
10
20
30
40
5C
60
30.99 ' 19.24
54.15 24.49
78.43
90.22
96.87
98.18
99.01
40.87
56.04
64.75
73.23
79.02
21.37
33.64
54.07
67.90
77.06
84.53
88.78
24.84
42.71
65.04
79.43
86.47
90.95
94.08
30.44
52.79
75.14
87.42
92.12
97.13
98.21
32.21
54.87
77.88
89.09
94.20
97.70
98.53
33.71
55.98
79.38
89,42
94.65
97.99
98.73
_I
Time
■"ater
5
44.54
10
72.76
20
93.71
30
98.31
40
1 99.53
Table IV.
Sodium Chloride at 25°.
rionoentratlon
_[_l'L„.^_i''_4
2 M
£1
-I
M/ii
*/*
21.80 28.93 35.44 42.23 44.65 46.91
36.85 50.48 I 60.97 ' 69.07 72.53 73.87
60.09 75.62 ' 84.93 90.30 91.90 03.51
75.73 87.85 93.45 97.56 97.95 98.57
85,90 , 93,91 1 97,28 , 98,75 99.26 , 99.69
I \ 1 I 1
29
Table V.
TlnK"
6
10
20
30
40
50
60
Calcium Chloride at 15".
Concentration
nrater
30.99
54.15
78.43
90.22
96.87
98.18
99.01
2.80
15.60
41.. 37
57.97
69.51
77.85
83.95
33.98
55.67
80.86
90.45
95.72
97.49
98.90
"•/
34.34
56.48
81.08
91.87
96.20
99.13
99.84
r-
34,01
56.58
81.19
92.04
96.96
98.25
99.19
Table VI.
Caloium Chloride at 25°.
Time
Oonoantratlon
Water
4 U
M
i:
44.54 20.18 48.36 48.61
10 72.76 49.93 75.24 76.23
20 93.71 78.20 | 93.76 | 94.52
30 98.31 92.23 99.81 97.94
40 99.53 , 96.16 100.00 99,3'
48.79
76.81
94.56
97.51
99.65
Table Yll.
M>gn«8lum Chloride at 15^.
30
Time
Concentration
,;ater
^ .
.,
../p 1
If/;
5
30.99
1.30
80,37
30.01
30.49
32.95
10
54.15
14.15
37.30
49.83
51.71
54.64
80
78.43
24.80
58.84
73.29
75.81
78.32
30
90,22
52.32
73.62
85.01
88.17
89.93
40
96.87
70.00
82.22
92.15
93.74
95.09
50
98.18
86.61
88.35
94.62
96.96
97.43
66
99.01
89.73
92.09
97.76
98.73
99.31
1
Table VII]
•
Ma
i^esium
Chloride at 25°.
1
Time '
Conoantratlon
7ater
4 M
2 M
H
t-1
43.09
5
44.54
1.80
29.52
40.01
41.60
10
72.76
25.36
56.29
69.41
70.89
72.08
20
93.71
73.30
79.97 ! 89.79
92.39
93.31
30
98.31
98.54
91.74
96.16
97.07
97.89
40
99,53
96,73
97.15
98. '.r
98.84
. ''' P
31
Table IX.
Barium Chloride
at 15°.
1
I
\
rime
r.oncentratlon
Wst«r
M
M/2
:•'. ^
5
30.99
30,51
31.49
34.48
10
54.15
50.56
54.98
56.88
20
78.43
74.68
78.60
81.05
30
90.22
86.83
90.30
91.09 ;
40
96,87
92.96
95.76
96.20
50
98.18
97.09
97.66
98.10
60
99.01
98.10
99.10
99.88
Table X.
Time
5
10
20
30
Barium Chloride at 25°.
■i
Water
44.54
72.76
93.71
98.31
99.53
Concentration
H
39.57
68.43
91.25
97.85
98.44
I - . —
m/2
42.04
73.23
92.00
97.99
98.72
M/4
46.89
75.90
93.94
98.19
98.85
32
Table XI.
Strontium Chloride at 15°,
Time
Conoentratl
on
Water
2 M
V.
'-
■•/4
5
30.99
80.80
27.44
32.33
35.^2
10
54.15
39.43
48.17
55.15
57.00
SO
78.43
63.31
74.88
78.86
80.89
30
90.22
76.46
87.19
90.56
92.07
40
96.87
84.93
94.40
95.89
96.30
"50
98.18
89.84
97.28
98.00
98.43
60
99.01
92.18
98.42
99.08
99.32
Time
6
10
20
30
40
Table XII.
Strontium Chloride at 25 °.
water
44.54
72.76
93.71
98.31
99.53
Coneentratlon
2 M
31.40
57.58
84.43
94.38
98.19
M
41.11
72.15
90.86
96.47
M/8
47.17
74.35
92.87
98.79
M/4
i_
47.89
76.01
94.37
98.65
99.64
33
Table JIII.
Time
b
10
20
30
50
60
Sodium Sulphate at 15°.
Yater
54.15
78.43
90.22
40 96.87
98.18
99.01
Concentration
M
M/2
m/4 ,
38.98
.37.16
65.30
60,57
86.15
83.74
94.07
93.27
96.64
96.43
98.32
98.13
99.65
98.97
Time
5
IC
20
30
40
Table XIY.
Sodium Sulphate at 25°.
Water
44.54
72.76
93.71
98.31
99.53
Concentration
M
61.51
87.56
98.16
99.38
99.83
m/b
54.62
82.84
96.95
99.22
99.71
m/4
50.52
79.32
96. C2
99.16
99.63
Table XV.
34
.V.agnealum Sulphate at 15°.
Tlma
Conoentratlon
Water
i
5
30.99
41.62
41.88
37.99
10
64.15
66.49
67.34
61.09
20
78.43
86.65
88.45
83.71
30
90.22
92.71
96.43
93.33
40
96.87
94.85
98.30
96.96
50
98.18
97.44
98.83
98.96
60
99.01
98.37
99.13
99.43
Ta
ble 7VI.
Maffneslum
Sulphate
at 25°.
1
J _,
f^onoentration ,
Ire
V/ater
:.V'3
5
44.54
56.61
50.98
55.23
10
72.76
84.26
78,72
80.46
20
93.71
05.49
95.82
95.99
30
98.31
96.48
98.17
98.29
^ (
99.53
98.13
99.71
no . ^"^
36
Table
XYII.
Potassium Nitrate at
15°.
1
1
Concentration
!
Tim*
Watar
••/2 ]
T//4
5
30.99
21.93
26. 5U
3u.5'J
J1.78
10
54.15
36.35
44.79
51.24
53.35
20
78.43
58.29
64.15
74.33
77.15
30
90,22
71.53
80.91
86.18
88.76
40
96.87
81.37
87.11
92.74
94.85
50
98.18
89.93
93.80
95.91
97,20
60
99.01
92.27
96.14
97.43
98,49
Table XVIII.
Potassium Nitrate at
25°.
1
Tine
Concentrati
on
Water
a M
U
M/2
M/4
5
44.64
34.44 35.56
37.03
41.64
10
72.76
51.90 54.37
65.93
69.71
ao
93.71
77.03 82.07
88.45
90.92
30
98.31 88.69
94.58
96.48
98.01
40
99.53 94.94 | 97.30
96,60
96.64 1
36
Tabic
XIX.
SodJ
Lum Nitrate at 15°.
Tlm«
Concantrat'
Ion
Water
»« iU
U
, li/E
-A
5
30.99
21.69
26.38
32.36
33.66
10
54.15
36.23
45.73
53.12
54.52
20
78.43
59.80
70.35
77.27
79.14
30.
90.22
73.51
81.95
87.93
89.11 ,
40
96.87
81.49
91.10
93.61
94.85
50
98.18
89.11
94.38
97.81
96.85
60
99.01
92.47
97.01
98.07
98.37
' Tlm«
5
10
20
30
40
Tabla XX.
Sodium Nitrate at 25 .
Conoantratlon
Water
44.54
72.76
93.71
98.31
99.63
2 H
33.14
36.80
54.72 63.93
I
80.20 j 86.68
91.87 95.77
9G.r' 7.89
41.07
69.01
90.58
97.06
98.24
41.64
69.94
90.92
97.54
98.83
37
DISCUSSION.
rher* Is one difficulty In th« study of this problem that
muBt first be pointed out, I.e., the use of a strong alkali so-
lution (half-normal NaOH) with which to titrate the acetic acid
forned. This necessarily introduces some error, slnee a dif-
ference of 0,1 cc. in reading the burette would make a differ-
ence of over 1 per cent. A more dilute solution of alkali
could not be used since too large a quantity of such a solution
would be required.
The rate of decomposition of the acetic anhydride is at
first very rapid, being almost complete at 25° in 5 minutes and
nearly three-quarters complete at the end of 10 minutes, then
gradually decreasing as the reaction approaches completion. In
this respect the reaction differs from similar ones studied,
such as the hydrolysis of esters, since in these cases the re-
actions are reversible. Temperature has a marked accelerating
Influence on the hydrolysis, the velocity of the reaction as a
whole, and the increase for succeeding Intervals of time, being
nraeh greater at 25° than at 15°.
All the salts studied, with the exception of sodium sul-
phate and perhaps also magnesium sulphate, have In the case of
the greater concentrations a retarding influence on the hydro-
lysis. This retardation diminishes as the salt solution be-
comes more and more dilute. ',7ith sodium sulphate solutions the
reverse is true, the more concentrated the solution the trreater
is the accelerating effect. This is also true to a certain
38
extent with ma^rneslum sulphate, although the effect la not as
pronounced.
In the oase of both ma^^eslum salts studied, ma/^nealum
chloride and ciairneslum sulphate. It was difficult to <?at clear,
clean-cut results. In titrating the acetic field with the alkali
in the presence of these salts a Tood end-point could not be
reached. Phe color of the Indicator, phenolphthaleln, appeared
to be maskei, especially In the more concentrated solutions.
All the non-hydrated salts studied have an hindering effect
on the hydrolysis. The amoxint of this hindrance under the same
conditions is practically the sane for the four salts studied,
there bein^ at no tine a variance of more than a few per cent.
Tith the most dilute solutions studied, ruarter-molar, the re-
sults for the decomposition are practically the same as for
pure water.
The hydrated salts, with the exception of ma^esium chlor-
ide, all ^ive results for the decomposition greater than those
of the non-hyrdated ones, while with the more dilute solutions
there is an appreciable acceleration of the hydrolysis of the
acetic anhydride over that due to pure water alone. Sodium
sulphate and magnesium sulphate at all concentrations studied
have a very narked acceleratinf? effect on the hydrolysis.
Greater concentrations of these salts were not used for the
reason that they do not mix with the anhydride at once on sim-
ple shaking. Calcium chloride, strontium chloride and barium
chloride also have an accelerating Influence on the hydrolysis
in the more dilute solution. L!agneslum chloride acts as do the
non-hydrated salts, having a retardlnf influence at all dilutions.
BIOGRAPHY.
Gerald Charles Connolly was born near rikesvllle, 3altlmore
County, I.!«ryland. on Au^at 7, 1890, He received hlg early edu-
cation in the public schools of Baltimore and receive! his A. 3,
decree from liock Hill College ('.'.aryland) . He then entered The
Johns Hopkins University as a graduate student in Chemistry
taking as his subordinates Physical Chemistry and Physics. Ho
received during the years 1914-15 and 1915-16 university schol-
arships*
39
■?^
