c. j. dobratz -- heat capacities of organic vapors.pdf

8/11/2019 C. J. Dobratz -- Heat Capacities of Organic Vapors.pdf

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June, 1941 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y 759

On the other hand, the low percentage of limiting-assay crudes

may merely strengthen the arguments of those who, in scout-

ing the predicted exhaustion of petroleum, contend th at the

production of crude oil

is

by no means a finished process, and

that new crudes are always being formed within the earth.

Acknowledgment

The author is indebted to his associates, particularly to

D.

W. Gould and 8.

.

Macuga, for supplying much of the

original data, and to the former for locating several of the

references cited. Both contributed helpful criticism and dis-

cussion of the theories advanced.

Literature

Cited

1) Birch, 8.

F.,

Dunstan, A. E.,

idler, F. A., Pim,

F.

B.,

and

Tait,

T., IND.NQ.CHEM.,1, 1079 (1939).

(2)Bur. of

Mines, Bull.

291 (1928), 01 (1937);

Tech. Papw

346

(1925);

Circ.

6014 (1926); Repts. Znveatigations 2202, 2235,

2290, 2293 (1921); 2322, 2364, 2416 (1922); 2595, 2608

(1924); 2807, 2808, 2824, 2846 (1927); 3253 (1934); 3346

(1937).

Candea, C., and Sauoiuc, L., Refiner

Natural

Gasoline Mfr. 18.

434 (1939).

11928>.

Cross,

Roy,

Kansas

City Testing Lab., BuU.

25, 284

et seq.

.----,

Egloff, Gustav, Lowry, C. D., Jr., and Schaad, R. E., J. Inst.

Petroleum Tech. 16, 133 (1930).

Egloff, Gustav,

Nelson,

E. F., and Morrell,

J.

C., IND. NO.

CHEM.,29, 555 (1937).

Hall,

F.

C.,

nd Nash,

A.

W., J .

Inst. Petroleum Tech.

24, 471

(1938).

Haslam,

R.

T.,nd Russell, R.

P.,

IND. NO.C H ~ M . ,2, 1030

(1930).

Houdry, E., Burt, W. F., Pew, A. F., Jr., and Peters, W.

A.

Jr.,

Proc. Am . Petroleum Inst. 111,19, 133 (1938).

Ipatieff,

V.

N.,

nd Egloff,

Gustav,

Natl . Petroleum News

May

16, 1935.

Ipatieff, V. N.,

and

Grosse, A.

V.,

IND.

NO.

CHEM.,28, 461

(1936).

Peterkin,

A. G.,

Bates, J.

R. and Broom,

H. P.,

preprint of

paper presented before Am. Petroleum

Inst.,

Nov. 17, 1939.

Rittman, W. .,Dutton, C.B.,nd Dean, E. W., Bur. of Mines,

Snelling,

W. O. U.

S. Patent 1,624,848 April 12, 1927).

BUZZ. 114, 8-64 (1916).

PRISINTIUDnder the tit le Petroleum Equilibria before the Division of

Petroleum Chemistry at the 100th Meeting of the Amerioan Chemioal

Society,

Detroit,

Mioh.

Heat Capacities of Organic Vapors

CARROLL J.

DOBRATZ,

University of Cincinnati, Cincinnati, Ohio

The method of calculating heat capacities of or-

ganic vapors from valence-bonding frequencies as

developed by Bennewitz and Rossner ( I ) has been

modified by the assumption of rotation within the

molecule so as to give more accurate results. Fre-

quencies have been assigned to valence bondings

with halogens, nitrogen, and sulfur which, when

used with those originally assigned to carbon, hy-

drogen, and oxygen bondings, enable the calcula-

tion of the heat capacities of the vapors of prac-

tically all organic compounds formed from these

elements. The values of the Einstein functions at

different temperatures corresponding to the as-

signed frequencies have been fitted to equations

of the form, - A BT

CTp,

so as to simplify

computations.

An

illustration shows the method

of calculating heat capacities, and comparison of

calculated and experimental data indicates that

an accuracy within 5 per cent can be expected in

most cases.

CCORDING to the work of Bennewits and Rossner

( I ) ,

the heat capacities of the vapors of organic com-

A pounds containing carbon, hydrogen, and oxygen can

be calculated by the following equation:

n - - Z*iC

1)

zqi

C v ) p

-

=

3R

+

where

R

= gas constant per mole

n = number of atoms in molecule

8 = number of valence bonds of ith type

yi, Csj = Einstein functions for a given bond

having

characteristic frequencies

v i and bi .

The Einstein function to be used in this case

is

the deriva-

tive of energy with respect to temperature:

where X = hv/lcT

h = Planck's constant

v = characteristic frequency

I = gas constant Eer molecule

T

= temperature, K.

e

=

base of natural logarithms

The Y frequencies were evaluated from Raman data, and

the 6 frequencies were determined empirically from experi-

mental data previously obtained

( 1 ) .

I n order to simplify

the computations, values of Einstein functions corresponding

to the given frequencies were listed at

40

intervals from

290 to 690 K. The expression,

which is derived from Berthelot's equation

of

state

I J ) ,

was

used to convert values of

(CJP

-

t o

the more commonly

used form (CJ,

.

Hea t capacities calculated by this method

were found to agree within 5 per cent with the experimental

data for a large number of compounds containing carbon,

hydrogen, and oxygen, mostly obtained a t the same tem-

perature, 410' K.

However, when the hea t capacities of propane, butane, and

pentane were calculated and compared with the experimental

data of Sage, Webster, and Lacey

( I @ ,

the differences were

much greater than 5 per cent at low temperatures (Table I).

This error is probably due to the fac t tha t this method takes

no account of rota tion within the molecule. The assumption


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760 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y

Vol.

33,

No.

6

of internal rotation has been frequently used, and recent

work on the ethane molecule tends t o bear out the validity

of the assumption. Several workers (4, 10 11,

17)

con-

ducted thorough investigations, and closest agreement of

theoretically calculated values of the thermodynamic func-

tions for ethane with experimental data was obtained by

assuming restricted internal rotation

of

the methyl groups,

using a restricting potential of approximately 3000 calories

per molecule. This restricting potent ial will allow virtually

free rotation at temperatures above 250 K. According

t o

Pitzer I 4 ) , the restricting potentials for the internal rota-

tions of larger hydrocarbon molecules are of approximately

the same magnitude, and therefore the assumption of free ro-

tation about carbon-carbon single bonds should be valid for

all temperatures above 300 K.

T A B L ~. HEATCAPACITIES

F

HYDROCARBONS

CP

at

1

atm.. cal./(mol.)(O

K.)

70'

F.

160' F. 250

F.

340'

F.

Hydroaarbon Souroem 294O

K.

344O

K.

394O

K.

444O

K.

Propane

a 1 7 . 8 1 9 . 2 2 0 . 8 2 2 . 3

b

1 6 . 7 1 8 . 5 2 0 . 9 2 3 . 3

1 6 .9 1 9 .0 2 1 .1 2 3 .4

:

1 5 .8 1 8 .1 2 0 .4 2 3 .0

n-Butane a

2 2 .9 2 4 .8 2 7 .0 2 9 .8

b

2 1 .4 2 4 .0 2 7 .2 3 0 .6

2 1 .3 2 4 .2 2 6 .9 2 9 .9

19.8 2 2 .9 2 6 .0 2 9 .4

n-Pentane a .

3 0 .6 3 3 .1 3 6 .8

b b . . . 2 9 . 8 3 3 . 5 3 7 . 3

0 ... 2 9 . 4 3 2 . 8 3 6 . 6

d

. . 2 7 . 8 3 2 . 0 3 5 . 9

a

a

= experimental values

of

Sage, Webster, and Lacey.(ib);

b =

calcu-

lated by method of Pitzer

14) ;

c = calculated from Equatlon

2;

d = calcu-

lated from Equation 1.

E Neopentane.

The average effect

of

the 6 vibrations as used in Equat ion 1

is

0 .5

calorie per degree of freedom for hydrocarbons a t

20 C. increases to approximately 1 calorie a t

150'

C., and

finally approaches a limiting value of 2 calories at higher

temperatures. The heat capacity due

t o

a rotat ional degree

of

freedom is

R / 2

or 1 calorie at all temperatures. Thus, at

137 C.

(410 K.),

the temperature a t which Bennewitz

and Rossner 1 ) obtained their experimental data, little

error would be produced by substituting a vibrational for a

rotational degree

of

freedom; but this error would appear a t

all other temperatures, since calculated heat capacities are

too low a t lower temperatures and too high at higher tem-

peratures. As Table I shows, heat capacities calculated by

Equation 1 agree within 5 per cent with experimental data

in the neighborhood of

410

K. but are too small at

the lower-temperatures. The assumption of free

rotation about carbon-carbon bonds in the mole-

cule with a corresponding decrease in the number

of vibrational degrees of freedom would tend to

correct this error.

It

is therefore suggested tha t

Equation 1be corrected

so

as to include internal

rotation.

THE changed equation

is

as follows:

a

=

number of bonds permitting free rotation,

In the same manner as for Equation 1

Berthelot's equation of state can be used to ob-

tain values of C , at finite pressures from the cal-

culated values of (Cv )p

0 .

C C

r similar (C-0 in ethers or esters)

The assumption of internal rotation has also been used in

the method of calculating heat capacities and other thermo-

dynamic functions as developed by Pitzer

(14) .

In this

method, which has been applied only t o hydrocarbons, heat

capacities are also determined by a summation of Einstein

functions corresponding to the vibrational frequencies

of

the

molecule. Approximately the same frequencies are used as

those employed by Bennewitz and Rossner

I ) ,

but the num-

ber of degrees of freedom to be assigned each vibrational

TABLE1. HEATCAPACITIES

F

ORGANICOMPOUNDS

CP

t

1

atm ., cal./(mol.)(O K.)

Eaperi- Calcd, by Calcd, by

Compound Temp., O

K.

mentala Equatlon 1 Equation 2

Acetone

Methyl ethyl ketor

Benzene

Toluene

Methylcyolohexane

ie

410

410

410

410

410

410

437

623

341

376

410

454

410

410

410

410

2 7 .5

2 7 .9

3 9 . 4

3 2 .9

3 3 . 3

1 9 .8

2 0 .9

2 8 . 2

2 0 . 1

2 1 .7

2 2 .5

2 3 .9

2 9 . 8

2 7 . 3

3 3 . 6

4 4 .5

2 7 .3

2 7 . 3

3 9 . 6

3 2 . 1

3 2 .1

2 0 . 0

2 1 . 0

2 7 . 1

1 9 . 8

2 1 . 3

2 2 . 8

2 4 . 6

2 8 . 8

2 7 . 3

3 3 . 6

4 4 . 8

2 7 .3

2 7 .3

3 9 . 9

3 2 .7

3 2 .7

2 0 .1

2 1 .1

2 6 .8

2 0 . 4

2 1 . 8

2 3 .2

2 4 .7

2 9 .3

2 7 . 3

3 3 .6

4 4 . 9

a

The experimental values at

410' K.

were obtained by Bennewitz and

Rossner

1 ) ;

values at other temperatures were taken from Landolt-BRrn-

stein

(fa).

frequency in a molecule is determined from a more elaborate

analysis of the structural formula, a separate analysis being

required for each different compound. Internal rotation

about carbon-carbon bonds in the molecule is assumed, re-

stricting potentials being such as t o allow practically free ro-

tation a t temperatures above 300 K. Heat capacities cal-

culated by this method for propane, butane, and pentane

are also listed in Table I and show closer agreement with

the experimental da ta than do those calculated by Equation 1.

Since Pitzer (14 ) also used approximately the same fre-

quencies as the ones assigned by Bennewitz and Rossner

I ) ,

those of the latt er were assumed

t o

be correct and were used

t o determine the Einstein functions for calculations with

Equation 2 .

As

Table I shows, heat capacities calculated

by means of Equation 2 for propane, butane, and pentane

agree as well with the experimental data as do those calcu-

lated by Pitzer's method and considerably better than do

values calculated by Equation

1.

Also, heat capacities

calculated for various other compounds (Table 11) show as

TABLE11.

BONDIXQ

REQUENCIES

ND

CONSTANTS T O EVALUATEINSTEIN

FUNCTIONS,

, = A

BT

CTa

Y Fre-

6 Fre-

quenoy uency,

Bonding Cm.-l

A

B

X 108 C X 10s %m.-1

A

B

x

103 c x

106

c-I

900 0.181 4 .664 -3 .338 260 1 .461 1 .730 -1 ,272

S

C B r

560 -0 .073 5 .158 -3 .59; 280 1 .242 2 .046 -1 ,501

C-Cl.

6 50 -0 .5 6 2 6 .3 8 5 -4 .4 9 5 350 1.023 2 .590 -1 ,874

c-s

C-C.

990

-1 ,0 9 0

6 .000 - 3 . 4 4 1

390 0.73 0 3.414 -2.577

C-0

1030

-1 ,1 7 3

6 .132 -3 .555 205 1 .461 1 .63 3

-1 .414

C-F

1050

-1 .128 5 .845 -3 .253

530 0 .011 5 .119 -3 .690

c c

1620

-0 ,4 3 2 1 .2 3 3 0 .9 3 6 846

-1 ,1 4 0 7 .2 5 4 -4 .93 6

C-0

1700

-0 ,3 2 4 0 ,7 2 4 1 .3 0 8

390 0 ,7 3 0 3 .41 4 -2 .5 7 7

C-N,

N-N

N L O

C L S

N L Q

C L N

S-H 2570

0 ,1 2 9 -1 .3 3 3

2.26 3 1050 -1,1 28

5.845 -3.2533

C-H

2920

0 ,2 2 9 -1 .2 2 4

1 658 1320 -0,938

3 .9 0 0 -1 .3 4 2

0-H 3420

0 .1 5 0 -0 .8 1 0

1.055 1150

-1 ,1 3 5 5 .3 6 3

- 2

740

N--IK


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June, 1941 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y 761

good as, or better agreement with experimental data than

do those calculated from Equation 1 by Bennewitz and

Rossner 1).

To illustrate the method of computat ion by Equation 2, a

sample calculation of the hea t capacity of acetone is pre-

sented. Quadratic equations of the form,

C =

A

BT CT2

as suggested by Fugassi and Rudy (6), re used to evaluate

the Einstein functions; the constants for the different bond-

ing frequencies are listed in Table

111.

6

C-H bonds, 2 C-C bonds, 1 C=O bond;

Zqi

= 9;

n = 1 0 ; a = 2

e

From the structural formula for acetone we find:

3n 6 a

Zqi 2

zqi

- 9

Summing up the equations for the different bonds,

For C-C:

2

Cv

2(--1.090

4

6.000

X

3.441

X

10-O 2

26.

For

0:

13 13

a =

-g- (+0.730 3.414

X lo-'

T 2.577 X lo-' Ta)

9

For C-H:

6 Cv

=

6(+0.229 1.224 X T 1.658

X

IO-

T )

26

2 6 Ca = -0.938 3.900 X

9 3

2 2

c6

9

9

$0.730 3.414 X

T 2.577

X

lo-' T*)

Cv

=

(-0.324 0.724 X

lo-'

T

1.308 X Tp)

T

-

1.342 X TI)

Z*JVi + 3n 6 a -

zqdc&

zqi

- 6.10 54.01 X lodaT - 18.43 X 10-6

T

aR

3R 7.95

2

C,),,o 1.85 54.01

X

T 18.43

X

10-6 T s

For acetone:

pc

= 47 atm.; T, = 508 K.

= 1.2 x 107/~8 at p = 1 atm.)

32

pe

Thus,

(CP) , - I

= 2.4

X

107/Ta 3.84

54.01

X

10-8

T

- 18.43

X

10-6 T

From this equation we obtain

C

for acetone a t

410

K. and

1 atmosphere pressure as 23.2 calories/(mol.)

(

K.) as com-

pared with the experimental value of 22.5 calories I),

a

difference of

3

per cent.

SINCE Equation 2 gives apparently accurate results for all

organic compounds containing carbon, hydrogen, and oxy-

gen, it should also be applicable to organic compounds con-

taining other elements, if the proper frequencies are

asso-

ciated with t he valence bondings.

The frequencies assigned by Bennewitz and Rossner 1)

should

also

be valid for the same bonds in compounds which

contain elements other than carbon, hydrogen, and oxygen.

Since these frequencies agree quite well with the average

values for the bondings

as

tabulated from Raman data by

Hibben

(8), t

was assumed that the average frequencies

as obtained from Raman data for organic bonds with halo-

gens, sulfur, and nitrogen could be used to evaluate Einstein

functions for use in Equat ion 2. The v and 6 requencies to be

associated with the various bonds are listed in Table 111.

Also given are equations

of

the form C,, A

BT

CTZ,

which can be used to evaluate the Einstein functions corre-

sponding to the given frequencies. These equations were

fitted

by

the method of averages to the values of the Einstein

functions at different temperatures as obtained from the

tabulated values in Landolt-Bornstein

(18)

. Deviations of

values of t he quadra tic form from true values are less than

1.5 per cent in the range 300' to 700 K .

The frequencies for carbon, hydrogen, and oxygen bondings

are those given by Bennewitz and Rossner 1). Those for

bonds involving chlorine, iodine, bromine, and sulfur are as

given by Hibben

8)

from Raman data. The frequencies for

the carbon-fluorine bonding were estimated by t he author

from Raman data

as

reported by Glockler and co-workers

(7).

Since there is little change in th e Raman spectra when a

nitrogen atom replaces a, carbon atom in an organic mole-

cule, carbon bonding frequencies were used instead of those

given for nitrogen bondings by Raman data. This decreases

the number of equations needed for computations and will

not materially affect the accuracy of t he results obtained.

There are not sufficient experimental heat-capacity data

available for nitrogen compounds

t o

justify more precise as-

signment of frequencies.

HEAT capacities for halogenated hydrocarbons were cal-

culated by means of Equation 2, using these frequencies,

and show good agreement with experimental data . In Table

IV these calculated heat capacities are listed along with ex-

perimental data and with the values calculated from a more

complete spectroscopic analysis of chlorinated and bromi-

nated methanes (16, 8). The differences are less than

5

per

cent except for the smaller molecules. This larger error is

probably due to the method of assigning frequencies.

In a molecule such as methyl chloride the deformation

vibrations could be associated with either the hydrogen-car-

bon-hydrogen or the hydrogen-carbon-chlorine bonding.

Because of t he differences in mass, neither of these bondings

will have

a

vibrational frequency equivalent to that of

a

carbon-carbon-chlorine bonding which is present in all higher

homologs. This is apparent in the Raman spectra, which

show no vibrations with a wave number below 712 om.-' for

methyl chloride, while the assumed 6 frequency for carbon-

chlorine is

330

em.-'. Hence the use

of

this

6

vibration for

methyl chloride introduces an error.

1

The equations

for

the carbon, hydrogen, and oxygen bondings were de-

rived by Fugasai and Rudy

8) rom

the table used by Bennewjtz

and

Rossner 1). They were checked by the author and found to be within the

stated limits

of

accuracy.

TABLE^

IV. H ~ A T

APACITIDSF

HALOQDNATDDYDRO-

C RBONS

C p ,

oal./(mol.)(o K.)

Compound Temp., 0 Exptl. Calod.

CHClzF

40 14.7'3 16.4

140 17.20 18.4

CChF

40 18.8' 19.3

140 21.2' 21.4

30.94 30.9

33.10 33.3

CClaF-CCLF

sq

140

CZHICL

110-220 22.7b

22.0

CsHsBr

28-116 17.5b 17.6

80-200

20.7b

20.0

CaHsCl

27-118 17.3, 18.1

120-230 18.7b 19.7

cclr

100 21.520 21.4

400 24.200 24.4

CHsCl

25 9.750 11.4

400 15.63' 16.5

CBrr

200 23. od 23.5

400 24.72d 24.6

CHaBri

200

17.18d 17.7

400 19.60d 19.5

a

Experimental values

from

Benning and McHarness

8 ) .

b Ex erimental values from International Critioal Tables 0) .

C

Cagulated by

Vold 18).

d Caloulated by Stevenson and Beach 18).


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762 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y Vol. 33, No.

6

In Table

V

heat capacities as calculated for mono- and

dimethylamines are compared with experimental values as

reported by Felsing and Jessen (6). The largest deviation is

about

7

per cent, but use

of

the higher frequencies for nitro-

gen bonds as given in Raman data, instead of the carbon

bonding frequencies, would increase this error.

This method

of calculation was not expected to give very good results for

ammonia, but as the values in Table

V

show, the differences

from accepted values are less than 5 per cent.

TABLE .

HEATCAPACITIESF NITROGENOMPOUNDS

C p

cal./(mol.)(o X.)

Compound Temp., C. Exptl. Calcd.

Monomethylamine

25

12.90 12.3

50 13.W

12.9

Dimethylamine

25

16.65 16.7

50

18.70

17.8

Ammonia

23 8.46 8.3

123 9.26 9. 2

223 10.06 10.1

323 10.7b 11.0

423

11.36 11.9

Ex erirnental

values

from Felsing and

Jessrr. 6 ) .

b C a h a t e d f r o m equations given by Bryant (3;.

Almost no gaseous heat capacity da ta are available for any

organic compounds of sulfw and therefore it was not possible

to obtain much of an experimental check on the frequencies

assumed for sulfur bondings. However, these frequencies

were chosen in the same manner as those for halogens and

should be of comparable accuracy. The one heat capacity

value found for an organic sulfur compound (that for diethyl

sulfide, 9) agreed within 2 per cent with the calculated

value. The calculated heat capacity

is

35.5 calories/(mol.)

K.)

and the experimental value, 36.2 calories for the range

120-223

C.

Literature Cited

1) Bennewitz and Rosaner,

Z .

physik. Chem. 39B, 126 (1938).

(2)

Benning and McHarness,

IND.

NG.CHEM.,

1, 912 (1939).

(3) Bryant, Ibid . , 25, 820 (1933).

(4)

Eucken and Part s, Z . physik. Chem. ZOB,

184 (1933).

(5) Felsing and Jessen,

J . Am.

Chem. SOC.,

5 ,

4418 (1933).

(6)

Fugassi and Rudy, IND.NG.CEEN.,

0,1029 (1938).

7) Glockler et al., J . Chem. Phys. 7, 278, 382, 669, 970 (1939);

8,

291 (1940).

(8) Hibben, Ram an Effect and I ts Chemical Applications, A. C.

S. Monograph 80, New York, Reinhold Pub. Corp.,1939.

(9) International Critical Tables,Vol. V p. 80, New York, MoGraw-

Hill Book Company,

1929.

(10)

Kemp and Pitzer, J . Am . Chem.Soc.

59, 276 (1937).

(11)

Kistiakowsky and Rice, J . Chem. Phys. , 7 ,

281 (1939).

(12) Landolt-Bornatein, Physikalisoh-chemisohe Tabellen,

Vol.

11,

(13) Partington and Shilling, Specific Heats of Gases,

p

40,

New

( 1 4 )

Pitzer,

J . Chem. Phys.

5, 47 3 (1937).

(15) Sage, Webster, and Laoey, IND.ENG.CHPM., 9, 1309 (1937)

(16)

Stevenson and Beach,

J.

Chem. Phys.

7, 25 (1939).

(17) Teller and Topley,

J .

Chem. Soc., 1935,885.

(18)

Vold, J . Am . Chem.Soc.

57,1192 (1935).

p.

1252,

Berlin, Julius Springer,

936.

York,

D.

Van Nostra nd Company,

1924.

.4BsTRAoTmD from a thesis presented in June, 1940,

t o

the Graduate School of

the University

of

Cincinnati in partial fulfillment of the requirements

for

the degree

of

master

of

science. The work

was

performed under th e direotion

of

E.

F. Farnau of

the

chemioal engineering faculty.

Reclamation of Stoddard Dry

J

Cleaning

Solvent

CHARLES

S. LOWE

AND

ADRIAN

C.

SMITH

National Association of Dyers and Cleaners

Silver Spring,

Md.

FTER repeated use dry cleaning solvent becomes con-

taminated with dirt, fat ty substances from perspiration

A and soap residues, grease, and oils

of

various kinds. The

insoluble soil consists of suspended matter such as lint, grit,

sand, dust, and water; the soluble impurities are usually

of an organic nature, consisting of mineral oils, fa tt y acids,

wool grease, unsaponified fats, dyes, and to some extent

various chemicals used in removing stains. The color of the

originally clear solvent, even though no soap or cleaning aid

is employed and no dyestuff is removed from the garments,

passes through shades of light yellow to dark reddish brown

as the poundage of clothes cleaned inoreases. Peroxides, de-

rived from the solvent itself by the action of light and from

oleic acid (the principal fatty acid constituent of dry cleaning

soaps), decompose and yield aldehydes, ketones, and low-

molecular-weight acids. Small amounts of such solvent re-

Present address, Pennsylvania Salt Manufacturing Company, Phila-

delphia, Penna.

The adsorptive properties of activated car-

bon, magnesiu m silicate, and activated

fullers earths toward fatty acids and sub-

stances associated with rancidity in dry

cleaning solvent have been studied in

laboratory and plant tests. Alkali ab-

sorption and mineral acid theories of dry

cleaning soap decomposition are presented

to

account for excessive accumulation

of fatty acids in Stoddard dry cleaning sol-

vent above that derived from the garments

thems elves. The effect of soap on the ad-

sorptive capacity of a powder is pointed o ut .

maining in the cleaned garments develop unpleasant rancid

odors. Eventually the solvent becomes so polluted as t o be

unfit for further application as a cleaning medium.

Before this condition is reached, it is necessary to clarify

the solvent by alkali treatment, by pressure filtration with

adsorbing powders, by vacuum distillation, or by a combina-

tion of these methods. It is generally held that distillation

will yield the purest solvent, removing insoluble soil and all

mineral oils, grease, and fatty acid residues having boiling

ranges over 410 F.which is the end point of Stoddard solvent.

c. j. dobratz -- heat capacities of organic vapors.pdf

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