I j88 DAVID R. MERRILL AND ELLIOT Q. ADAMS.

[CONTRIBUTION FROM THE CHEMICAL LABORATORY OF WE UNIV~~RSITY OF CALIFORNIA. ]

TYE HYDROLYSIS OF ACETANILIDE. B Y DAVID R. MBRRILL A N D ELLIOT Q ADAMS

Received June 5 , 1917.

The acetylation of aniline has been studied by Menschutkin' and by Goldschmidt and Waclis,2 but their experiments were made for the most part in glacial acetic acid or in aniline, hence the laws of aqueous solutions could not be expected to hold. The hydrolysis of acetanilide in alkaline solution has been studied by Davis." While the hydrolysis of other anilides and acetyl compounds in acid solution has been studied, little has been done on the hydrolysis of acetanilide, which is probably the most important. For these reasons it was thought worth while to measure the rates of re- action and equilibria concerned in a fairly dilute acid solution.

n

Air 4 1 d t 4 -

U

4 Experimental Work.

-The progress of the reaction was followed by determining the total acidity by titration. In the acetylation experi- ments solutions were made up of aniline and

--- acetic acid,, and mea- --- sured out with a pipet

into glass tubes which

--- - - -

were then exhausted to a b o u t 25 mm. a n d sealed. The tubes were placed in a thermostat at IOO', heated the re- quisite length of time, opened and the contents washed out and titrated. In the hydrolysis with acetic acid alone, the same amount of acetan- ilide was weighed out into each tube and a pipetful of acid added,

Fig. I . since acetanilide is not 1 Menschutkin. Chem. Zentr., 1906, I, 5 5 1 . 2 H. Goldschmidt and C . Wachs, 2. phys. Cltem., 24,353 (1897) 3 0 M C. Davis, J Chem. Soc., 95, 1397 (1909)

THE HYDROLYSIS OF ACETANILIDE. 1589

soluble enough at room temperature to permit the making up of a solution of suitable concentration.

In the case of the hydrolysis in the presence of hydrochloric acid the reaction proceeds so much faster that it was found more convenient to use the apparatus shown in Fig. I , The water and acetanilide were placed in the flask and the hydrochloric acid in the pipet A. After the whole had come to the temperature of the bath and the acetanilide was all dis- solved, the acid was blown over into the acetanilide solution and thoroughly mixed. Then by turning the stopcocks in the proper sequence a pipetful of the solutiop could be blown out for analysis. By always running the solution into the same amount of cold water and titrating immediately, fair uniformity of the conditions was obtained. Because of the method of calculation used, the actual time from the beginning was not so im- portant as uniformity in the measurement of the time of drawing the sample.

Method of Calculation.-We shall consider first the simplest case, namely, the hydrolysis in the presence of hydrochloric acid. Since at the concentration of acid present throughout the experiments the ionization of acetic acid and the “hydrolysis” of aniline hydrochloride are largely repressed-while the concentration of acetanilide hydrochloride is negli- gible-the ionic equation for the reaction studied is

As the concentration of the water (more correctly its activity) does not change materially, the reaction may be expected to have the character of a bimolecular reaction, the rate being proportional to the product of the concentrations of acetanilide and hydrogen ion. (Under the ex- perimental conditions the reverse reaction is entirely negligible.) The equations for the progress of such a reaction are

CsHsNHCOCH3 + HzO + H + --j C6H5”3+ + CH1COOH.

dx = k2(a - x)(b - x)dt b--x b 2.303 log - = (b - a)& + 2.303 log - a - x a

where a, b and x are, respectively,’ the acid and acetanilide originally present, and the amount of each which has been used up a t time t . The correction that must be made for the hydrogen ion furnished by the acetic acid and by the aniline ion is as follows: The ionization constant of acetic acid2 a t 100’ is 1.11 X IO-^, that of aniline ion3 3.1 X IO-^, so that a mixture of equivalent amounts of the two in the presence of excess mineral acid will give as much hydrogen ion as one equivalent of an acid of con- stant 3.2 X IO-^.

The notation differs from that used later in the paper in the calculation of the experiments in acetic acid solution.

If y represents the hydrogen ion so produced

* A. A. Noyes, THIS JOURNAL, 30,335 (1908). a Determined colorimetrically, using methyl orange as an indicator.

1 5 90 DAVID R. MERRILL AND ELLIOT Q. ADAMS,

3.2 X IO-~(HA) = (A-)(H+) 3.2 X I o u 4 (x -- y ) = y (n - x + y )

or 3.2 X IO “(x-y)

from which y may readily be calculated by approximation. Since the rate of the reaction is directly proportional to the hydrogeu-ion concentration, correction for the extra hydrogen ion y may be made by increasing all

measured time intervals in the ratio of I + ___ to I . In Fig. 2, log - G - X a - x

has been plotted against the “effe,ctive time,” and from the slope m of the

Y = a - x + y

b--x

C) I 2

Effective time in horns. Fig. 2.

3

straight lines on these plots, the constant kt of the hydrolysis reaction has been calculated by the equation

2.303 m kz = -. b - a

The data and calculations are given in Tables 2 and 3. When the reaction takes place in acetic acid solution the speeds of the

the two directions are comparable and the hydrogen ion from salt and acetic acid, instead of being a small addition to that

already present, constitutes the entire hydrogen-ion concentration of the solution. In addition the aniline and acetic acid are no longer present in equivalent amounts. The equation for the progress of the reaction

CHaCOOH + CsHsNHs+ CsHaNHCOCHs + HzO + H+,

THE HYDROLYSIS OF ACETANILIDE. 1591

since the activity of the water is substantially constant, takes the form

where x represents the concentration of acetanilide. If a and b be, re- spectively, the total acetic acid and aniline, each including acetanilide, and x the acetanilide present at time t, this equation can be put, according to the direction in which the reaction is proceeding, into the forms

dx (a - x ) (b - x )

d~ = [~~(CH&OOH)(CBHSNH~+) - k2~(H+)]dt (1)

- -

(112) ki (CHaCOOH) (CsHs"3') dt. 1 - dx = kz [(H+) - -

X k2 X

TABLE I. CcHJVHz + CH&!OOH = CeHsNHCOCH3 + H20.

No. Dir.1 ZAn. ZAc. (HCl). Remarks. No. Dir. CAn. ZAc. (HCI). ,Remarks. I + 0.1 0 .2 ... 2 - 0.1 0.5 . . . 3 - 0 . 1 1.0 . . . 4 ---t 0.3 0 . 5 . . . 5 --+ 0 .3 1.0 . . . 6 4 0 . 3 1 . 5 . . . 7 f- 0 .3 0 . 7 . . . 8 -0.3 1.0 . . . 9 .C- 0.3 1.5 ...

IO 0.3 0.7 . . .

AnilinetoodiluteII + 0.3 1.0 . . . Same 12 + 0.3 1.5 . . .

Acid too dilute 14 0.3 1 . 0 . . . SeeTable4 15 0.3 1.5 . . . SeeTable4 16 C- 0.2 0.1

See Table 5 17 C- 0.2 0.1 See Table 5 1 9 ~ t 0.2 0.1

See Table 5 20 C- 0.2 0.1

SeeTable4 ZI + 0.2 0 . 1

Same I3 0.3 0 . 7 . . .

Duplicate of 5 Duplicate of 6 See Table 6 See Table 6 See Table 6 See Table 2

See Table 2

See Table 2

See Table 2

See Table 2

The mass-law equations for acetic acid and aniline ion take the form

The arrow shown is to be substituted for the = sign in the equation. Series 18 was lost because the acetanilide was not all in solution.

* This is not the usual KB of aniline, but its reciprocal multiplied by the dissocia- As no data were available for this temperature, we have tion constant of water.

determined K'B at 100' for aniline to be 3.1 X IO-(.

DAVID R. MERRILI, AND ELLIOT Q. ADAMS.

. . . . . . . . . . . . . . . . . . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . g 6 L ei G 4 .d b m

THE HYDROLYSIS OF ACETANILIDE. I593

d.

00 m ?

s m

N

00

m m

f

?

s h N m

N

CQ vi

..

v!

3

s

a

0, m

m N ?

e m

0

m m m

m \o rg

?

\o : s ! m

H

'9 v) m

s h

m

rg

m

?

'9 m

I- n a s h

m i!

z 'on 8 5 m

\o

N m c?

I- n

cl.

2 s m

0

h m

s 8 d.

N ? s

h '9 s s v)

m

m

h m

F!

v!

s 2

z x

d.

d . w $jp??o"% ? n p ! o o o o 0 d o o

m s 0

m v)

? 0

w CI

8 0

'9 0

d m

s t? m

0,

m h M

0

W d ? 0

0,

3 0

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8 I?

w

h m

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h h ? 0

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m 0 . n . 2

m n : v ) .

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s . m - m

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d ! ? .

. . . . . .

. . . . * . . . . . . .

h s I

0

5

m h P

2 c

b

I594 DAVID R. MERRILL A N D ELLIOT Q. ADAMS.

TABLE 3. Slope kt, Barometer.

Series No. recip. hr. I/mol X hr. Temp. obs. in. (of Hg). Temp. calc.

20. . . . . . . . . . . . . . . . . . . . 0,349 8.10 99. j ' 29.81 99.90' 2 1 . .................... .0 .352 8.17 99.5' 29.45 99.55'

Mean . . . . . . . . . . . . . . . . . . . 8.30 99.83' . . . 99.83O . . . . . . . . . . . . . . . . . . . . I00.00' Reduced . . 8 . 4 3

Series 16 and I 7 were performed in an oil thermostat and gave for kz, 7.46 and 7.92, respectively. An investigation into the cause of this discrepancy showed that the thermostat regulator was not operating satisfactorily. The subsequent experiments were performed in a boiling water bath and the result reduced to IOO', assuming a 2l/a-fold increase for IO'.

Substituting in Equations (111) and (112), since kl/k2 = K, the equilibrium constant for the acetylation of aniline (determined in separate experi- ments, see Table 6) ,

r9 . . . . . . . . . . . . . . . . . . . . 0 . 3 i 2 5 8.64 1 0 0 . 3 ~ 29.97 100.05°

or

Now from (IVa), (IVb), (Vu), (Vb), since in the absence of other elec- trolytes

(H+) = (CH&!OO-) - (CeH&Ha+), (VIII)

a - x U - - x

2

whence (H+) may be calculated by approxifiation. Then by plotting,

respectively, log (:E:) - and -log x against the "effective" values for the

time, obtained by multiplying every measured time interval by the mean value of the respective expression in square brackets, straight lines should

THE HYDROLYSIS OF ACETANILIDE. I595

be obtained from whose slopes, ml, mz, the reaction velocity constants can be calculated, as in the earlier case of hydrolysis by hydrochloric acid, by means of the equations

The results are shown plotted in Pigs. 3 and 4; the data and calculations will be found in Tables 4, 5 and 6. The concordance is .not all that could be desired, but since the equilibrium experiments indicated side

TABLE 4. Time.

Sesies No. Hrs. 6-2

3

4

5

6

7

5-1

2

3

4

5

6

7

1-1

2

3

4

5

6

6.5

IS .o

28.5

48.75

72 .o

98 .o

I .g2

7 e42

18.92

29.42

49.67

72 *92

98.92

26

47

97

I93

361

697

X Log mol/l. (a-s)/(bs)

0.01476

0.02970

0.04877

0.06955

0,0899

0.1107

0.00329

0.00776

0.02061

0.02ggo

0.04432

0 .os80

0.0717

0,01778

0.02381

0.04504

0 -0729

0.1027

0.1265

0.720

0.739

0.768

0.801

0.836

0.879

0.530

0.535

0 . 5 5 0

0.560

0.580

0.598

0.618

0.386

0.392

0.414

0.446

0 * 490

0.531

IOC(H+) mol/l.

1.598

I .646

I .712

1.798

I .895

2 .OIS

I .22I

1.229

1.255

1.274

1.309

1.343

1.384

0.998

I .006

1,037

I .086

1.153

I .222

Factor. ( I.

0.312

0 - 309 0.302

0.283

0.276

0.254

0.2579

0.2575

0.2528

0.2490

0.24.4.4

0.2370

0.2274

0.2086

0.2059

0. I933

0.1696

0.1266

0.0673

3.56

3.21

6.04

6.60

6.88

I .42

2.93

2.63

4.99

5.60

6.04

4.36

9.98

17.41

24.89

32 e58

. . .

3.56

6.77

12.81

19.41

26.29

...

1.42

4.35

6.98

11.97

17.57

23.61

. . .

4.36

14.34

31.75

56.64

89.22

1596

Series No. 9- 1

2

3

4

5

8-1

2

3

4

5

7-1

2

3

4

5

DAVID R. MERRILL AND ELLIOT Q. ADAMS.

TABLE 5 . X

md/l.

0 . 2 3 2 7

0.2189

0 . 2 0 2 2

0.1956

0.1867

0.2407

0.2267

0.2092

0.1958

0.1860

0,2546

C.2373

0.2181

0.2039

0.1912

ZC&COIH ZMsNHz Series G b No. mol/l. mol/l. 6 1.437 0.2860 5 0.963 0.2862 IO 0.672 0.2870

1 f l o g x. 0.36 j 7

0.3402

0.3058

0.2914

0.2711

0.3815

0.3554

0.3206

0.2918

0.2695

0.4059

0 ' 3753

0.3387

0.3094

0.2815

IOa(H+) mol/l. 4.188

3.604

3.110

2.968

2.800

3.246

2.746

2.359

2.1.50

2.030

2.821

2.196

I ,832

I .662

I ,547

lo'[ I lO'AtX[ Imem low mol/l. mol X hr./l. mol X hr./l. 3.902

3.268

2.716

2 3 5 2

2.354

3 ' 105

2.573

2 . 1 5 1

I .916

I .776

2.763

2.117

1.730

1.543

1.4'4

TABLE 6. Slope 2.303m/a-b

Tn ki ka recip. hr. I/mol X tu. I/mol X hr. 0.00604 0.01209 0.00374 0.01272 O.OOIg74 0.01180

I/mol x hr. 9 1.429 0.2860 2.65

7 0.670 0.2870 3.65 8 0.954 0.2862 3.09

I 5 1.442 0.2860 14 0.963 0.2862 13 0.67~ 0.2870 Extrapolated to 0 conc. of CHaC02H

86. I

143.7

1 2 1 . 2

181.5

68.1

113.3

93.6

92.3

58.6

100.1

68.6

109.3

K.

. . .

86.1

229.8

351 .o

532.5

. . .

68.1

181.4

275 .o

367.3

. . ,

58.6

158.7

227.3

336.6

6.10 7.12 8.41

7.92 X IO-' 8.78 X IO-* 9.32 x IO-' 10.54 x IO-*

reactimsssuch as the formation of diphenylamine (which was in fact shown qualitatively to be fornied)-a kigh degree of concordance could hardly be expected. There is the possibility that in addition to the

All the tubes used in the quantitative work developed a more or less pronounced magenta color, even though they were exhausted down to 25 mm. before sealing. Subse-

THE HYDROLYSIS OF ACETANILIDE. I597

catalytic hydrolysis of acetanilide by hydrogen ion and by hydroxyl ion CeHsNHCOCHa + H+ + H2O = CeHsNH3' + HOCOCHa

CsH6NHCOCHa + OH- = CeHsNHz + OCOCHs- direct hydrolysis by water may occur

a t a rate which is negligible in solutions containing an appreciable amount of strong acid or alkali and yet large enough to constitute the greater part of the hydrolysis in neutral solution.

CaHsNHCOCHa + H2O = CeHsNHz + HOCOCHs

0 . 9

0.8

H H 0 . 7

0.6

OM 0.5 cl 0 .4

0 . 3

3 i . 3 rl

- 1.2

0 IO 20 30 40 o 0.01 0.02 0.03 0.04 Effective time in hours. Hours X mols per liter.

Fig. 3. Fig. 4.

The results of the experiments on the hydrolysis in the presence of hydrochloric acid agree within the constancy of the temperature, about 0.2', corresponding to 2% on the value of k2, so that the mean value, reduced to IOO', may be taken as 8.4 0.1. This result is to be compared with the corresponding constant found by Davis' for the hydrolysis by alkali, namely 3.2 (at the boiling point of tenth molal solution). The most probable value of the acetylation con- stant K I may be found by extrapolating2 the equilibrium constant K to zero concentration of acetic acid, and multiplying by kz. The values thus obtained are K = 10.5 X IO-^, kl = 0.89 X IO-^.

summary. I. The hydrolysis of acetanilide by hydrochloric acid proceeds in accor-

dance with the equation CsHsNHCOCHs + H + + H20 = CsHsNH3' + HOCOCH3.

2. The velocity constant a t 100' for the hydrolysis of acetanilide by hydrochloric acid is 8.4 f 0.1 liters3 per mol per hour. (Davis found 3.2 for hydrolysis by alkali.) quent experiment has shown that continued boiling out under this pressure entirely prevents the development of color, but to do so without change in concentration seemed impracticable.

Davis, LOC. cit . Both constants are in liters per mol per hour, * Direct extrapolation of the values of K1 gives a result consistent with this, but the

3 Concentrations are stated in mols per cubic decimeters a t IOO". curvature of the graph makes extrapolation inconclusive.

1598 FRANCIS P. VENABLE AND JAMES M. BELL.

3. The moat probable value (obtained by extrapolation to zero concen- tration of acetic acid) of the equilibrium constant (omitting the concen- tration of the water) for the reaction as written above is 950 (= I/IO.S X IO-*)); this corresponds to a velocity constant for the acetylation of 0.0089 liters per mol per hour.

In conclusion we wish to express our gratitude to Prof. H. C. Biddle for his many helpful suggestions.

BBPSELGY, CALfPORNIA.

[CONTRIBUTION FROM THE DEPARTMENT OF CHEMISTRY OF THE UNIVERSITY OF NORTH CAROLIN-4. ]

THE ATOMIC WEIGHT OF ZIRCONIUM. BY FRANCIS P. VBNABLE AND JAHBS M. BELL.

Received June 8, 1917.

As the results of our work involve a considerable change in the atomic weight of zirconium, it is well to give a brief review of the previous deter- minations of this constant.

TABLE 1.’ Previous Determinations of the Atomic Weight of Zirconium.

Number of Atomic Date Chemist. Ratio analyses. weight.

. 1825 Berzelius* Zr(S04)~:Zr02 6 89 46

1889 Bailey4 Zr(SO&: ZrOz 8 9 0 . 6 5 4 . . . . , . . . . 1881 Weibul13 Zr(SeO&: ZrOz 5 90 * 7 9

, 1860 Marignac5 KzZrF6:K2S04 4 90.03

1844 Hermann’ ZrClr. ? I aa . 6 4 1844 Hermann’ zZrOCiz.gH20: ? 2 89.98

1881 Weibul13 Zr (SO& : ZrOz 7 8 9 . 5 4

6 . . . . . . . . 1860 Marignac6 KZZrF6:ZrOz 4 9 1 . 5 4 . 1898 Venable& ZrOClZ.3HzO : ZrOs IO 9 0 . 8 1

We believe that little dependence is to be placed upon any of these figures. It is now well known that this substance cannot be freed from water without losing some sulfur trioxide. In addition it is difficult to prepare pure, especially avoiding the products of hydrolysis.

The series of determinations in No. 7 was carried out by one of us in 1898 and a t the time of publication several sources of error were painted out and a promise given of further work on the subject. Circumstances have prevented the fulfillment of the promise until now. The most serious

In the first three, zirconium sulfate is used.

‘ Clark, “Constants of Nature,” p. 287 (1910) Ann. Physik., [ I ] 4, 126 (1825).

3 Lund. Arsschrijt, 18 (1881-2). 4 Chem. News, 60, 6, 17, 32 (1889). 6 Ann. chim. p h y s , [3] 60, 2 7 0 (1860). 6 THIS JOURNAL, 20, I 19 (1898). 7 J prakt. Chem., 31, 75 (1844).