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I. Isotopic studies: A. A redetermination of the protium-deuterium ratio in normal water. B. The microdetermination of the density of small quantities of water by the balanced drop method. II. A study of the properties of lithium glass electrodes

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NAME AND ADDRESS
DATE
FOR THWES TERN UNIVERSITY
I. ISOTOPIC STUDIES
A.
A REDE TERMINA TION OP THE PROTIUM-DEUTERIUM
RATIO IN NORMAL RATER
B.
THE MICRODETERMINATION OP THE DENSITY OP SMALL
QUANTITIES OP WATER BY THE BALANCED DROP METHOD
II. A STUDY OP THE PROPERTIES OP
LITHIUM GLASS ELECTRODES
A DISSERTATION
SUBMITTED TO THE GRADUATE SCHOOL
IN PARTIAL FULFILLMENT OP THE REQUIREMENTS
for the decree
DOCTOR OP PHILOSOPHY
DEPARTMENT OP CHEMISTRY
BY
JAMES LAWRENCE GABBARD
EVANSTON, ILLINOIS
MAY* i&99
IL c
n e,
I9
*^-0
ProQuest Number: 10101426
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Table of Contents
I* Isotopie Studies
A*
A Redetermination of the Protium-Deuterium Ratio
in Normal Wa^-er
Page
I.
Review of Literature
1
II.
Statement of Problem
6
III. Experimental Methods*
6
1. Preparation of Deuterium-free Water
6
2. Preparation of Water Composed of
Deuterium-free Hydrosen and Atmospheric
Oxygen
8
5* Preparation of Rinsing Water
9
4. The Purification and Density
Measurements of the Water1
9
5. Preparation of Light water from a
Mixture of Light Waters
15
IV* Discussion of Data and Comparison
with Previous Results
V.
Calculations of Some Isotopic Exchange
Equilibrium Constants
16
25
VI . Summary
29
VII .Bibl io graphy
31
B.
Microdetermination of the Density of Small Quantities
of Water by the Balanced Drop Method
Page
I.
Review of Literature
1
II.
Statement of Problem
7
III* Materials and Experimental Procedure
9
1.
Preparation of Cyclohexyl Chloride
9*
2*
Measurements of the DensityTemperature Coefficient of the
Oil Saturated with Standard Water
10
3*
Some Preliminary Experiments
13
4.
Density Measurements
19
5*
Solubility of the Oil in Water
19
6.
Solubility of Water in the Oil
20
IV.
Discussion of Data
22
V.
Summary
29
VI.
Bibliography
3d
II
A Study of the Properties of the Lithium Class
Electrode
I.
Review of Literature
1
II,
Statement of Problem
3
Ill* Some Preliminary Experiments
IV*
1.
Apparatus and Materials
4
2*
Experimental Procedure
8
Materials and Experimental Procedure
for the Study of the Lithium
Glass Electrode
10
1*
Preparation of the Electrode
11
2.
Resistance of Lithium Glass
Electrodes
IS
The Vacuum Tube Potentiometer
for the High Resistant Lithium
Electrode Measurements
13
The Asymmetric Potentials of the
Lithium Glass Electrodes
17
The Sensitivity of the lithium
Glass Electrode to pH Change
18
A Comparison of the Behavior of
Lithium Glass Electrodes with
the Behavior of Sodium Glass
Electrodes of Approximately the
Same Resistance in the Same
Solutions
20
3*
4*
3,
6.
V*
4
Conclusions
VI* Bibliography
24
Tables
Isotopic Studies
A' Redetermination of the Protium- Deuterium Ratio
in normal
Water
Pa ere
Table I,
U:
II.
Data Obtained in the
Preparation of Deuteriumfree Water
7a
Data Obtained in the
Preparation of Water made from
Deuterium-free Hydrogen and
Atmospheric Oxygen
7a
!t
III.
An Illustration of Data as Taken 14a
in
IV.
The Increase of the Equilibrium
Temperature of the Float in
Standard Water with Time
15a
Data for the Density of the
Various Waters
17a
Difference in Density of water
Containing Formal Oxygen and
Atmospheric Oxygen
18a
V.
VI.
VII.
Density Differences Between
Deuterium-free Water and Formal
Water from Different Investigations
18a
u-
VIII.
Force Constants of DgO Molecule
28a
Microdetermination of the Density of Small Quantities
Of Water by the Balanced Drop Method
Page
Table I.
Density-Temperature Coefficient
Data
12a
Page
Table IX,
11
"
11
11
"
u
III,
IT,
V.
TI,
11
VII
Bata Illustrating a Typical
Measurement
15a
Equilibrium Temperatures of Water
N and Standard Light Water G in
Oil in Preliminary Experiments
16a
Equilibrium Temperatures of the
Standard and Light Standard
Waters
19a
Bata for the Bensity of the
Tarious Waters
19b
Solubility of the Oil in Water
at Approximately 25° C,
20a
Solubility of Water in the Oil
at Approximately 25° C,
20a
Bata Illustrating a Typical
Measurement
24a
II,
A Study of the Properties of the Lithium Glass
Electrode
Table I,
w
“
II,
III,
IT,
Potential Measurements of
Hydrogen-Calomel Cell B and
Glass-Calomel Cell C
10b
Errors of the Glass Electrode
at 30° C.
10c
Resistance in Megohms of
Lithium Glass Electrodes
14a
Vacuum Tube Potentiometer
Measurements on a Standard Cell
with Known Resistances Inserted
17a
Page
Table V.
Asymmetry Potentials of Lithium
Glass Electrodes
18a
11
VI*
Saturated Calomel-Glass Electrode
Cell
at Approximately
250 C.
19a
"
VII.
Bata Comparing the Behavior of
Lithium Glass Electrodes with
Sodium Glass Electrodes of
High Resistances
”
VII.
21a
Saturated Calomel-Glass Electrode
Measurements
21a
Graphs
Microdetermination of the Density of Small Quantities
of Water by the Balanced Drop Method
Fig.
"
u
'*
la*
lb*
2.
3*
Abs'olute Density of the
Saturated Oil and Water as a
Function of Temperature
13a
Net Differences in the Densities
of the Saturated Oil and Water
13a
Equilibrium Temperature of G
Water in the Saturated Oil
15b
Equilibrium Temperature of B
Water in the Saturated Oil
24b
A Study of the Properties of the Lithium Glass'
Electrode
Fig*
1*
Potentiometer Circuit
4a
Comparison Cell
Error of the Glass Electrode at
C« Compared with the Data of
Dole and Wiener at 25° and50 C
Input Circuit for Measuring
the Resistance of Glass
Electrodes
Input Circuit of Potentiometer
for Measuring the Potential of
Very High Resistance Electrodes
Acknowle dgment
The writer Is indebted to the
University of Kentucky for a sabbatical leave of
absence and to Northwestern University for a
Tutorial Fellowship both of which made it possible
for him to continue these investigations*
He wishes
to give sincere thanks to Mrs. Mary Bradley Gabbard,
his wife, and to Ur. Franklin E. Tuttle, Professor
Emeritus, of the Chemistry department, University
of Kentucky, for their interest and aid in the
preparation of this manuscript.
He wishes
particularly to express his obligations to Ur. Malcolm
'
\
Dole, Assistant Professor of Chemistry, Northwestern
University, who directed these investigations, for his
guidance and assistance.
Further, thanks are eiven
to the many individuals who contributed in many ways
to the progress of these investigations.
Cl)
A Redetermination of the
Deuterium- Protium Ratio in Uormal Water
Review of Literature
Even before the discovery of the isotope
of hydro'gen of atomic mass 2, by Urey and his co/«» v
[g j
workers,vx; Birge and Menzel
had predicted that
the discrepancy between the chemical atomic weight
of hydrogen of 1*00777*0.00002 and Aston1s spectrograph
value reduced to a chemdcal value of 1.00756*0.000015
would be removed by an isotope of hydrogen of mass
2 with an abundance ratio of H/D in ordinary water
of 4500*
Prom intensity observations made by Urey
and co-workers upon the spectrograph lines of
hydrogen of mass 2,it was estimated that the
abundance ratio of H/D should be 4,000.
Upon examin­
ing a sample of commercial electrolytic hydrogen
(5 )
with his mass spectrograph, Bleakney' ' reported
this value to be 30,000,which was confirmed at the
1.
Urey, Brickwedde and Murphy, Fhvs.rev.,39,164(1932);
"
n
40,1 (1939)
2.Birge and Menzel, Phys. Rev.,37,1669 (1931)
3. Bleakney, Phys. Rev., 41,32 (1931)
(2)
time by other investigator.*^' However, Bleakney
/c\
and GouId
later reported a value of 5000 for the
ratio in rain water and assumed that their previous
value of 30,000 was due to electrolytic separation
of the isotopes by electrolysis.
By electrolysing
ordinary water in electrolytic cells of known
separation efficiency, Lewis and Macdonald
(6 )
estimated the ratio to be 6500, and many other
investigators have reported various values for
this ratio' (Table VIi#se<e belot P«1Qbl) .
rnh© wide
discrepancy between these reported values is due to
the failure of some of the investigators to correct
for the fractionation of the oxygen isotopes in the
preparation of their water samples, and to the
employment of different methods by others in correct­
ing for this error.
Johnttcrr(^^one of the first to correct
for the oxygen isotope fractionation, reduced 50
liters of ordinary water to 45 cc. by five successive
4.
5.
6.
7.
Hardy, Baker and Dennison, Phys. Rev., 42,229 (1932)
Tate and Smith, Phys. Rev.,43,672 (1933)
Bleakney and Gould, Phys. Rev.,44, 265 (1933)
Lewis and Macdonald, J* Chem. Phys.,1,341 (1933)
Johhston J. Am. Chem. Soc., 57, 4B^ (1935)
(3)
stages of fractional electrolysis between iron
electrodes, which were known to have an electrolytic
separation factor of eight towards the isotopes of
hydrogen#
Twenty-five percent of the electrolyte
was decomposed in each stage of the electrolysis,
and the water produced by burning the electrolytic
gases from a pyrex jet served as the electrolyte in
the next stage#
The density of a sample of the
original water and a sample from each of the
electrolytic stages was determined by the submerged
float method and was plotted as a function of the
electrolytic stage#
After the third stage of
electrolysis, the decrease in density per stage
was constant.
Since a cell with a separation factor
of eight for the isotopes of hydrogen would produce
deuterium-free water after the third stage of
electrolysis, it was assumed that the decrease in
density of the water in the succeeding stages was
due to the fractionation of the isotopes of oxygen.
By extrapolating over the straight line portion of
the curve the density of deuterium-free water was
found to be -18#3 Y/The symbol Y represents in p.p.m.
(parts per million) the density of the deuterium-free
water less the density of normal water,from which
value the H/D ratio was found to be 5800*
(Q \
Morita and Titani' 7 corrected for the
oxygen isotopes in two ways,
first, they combined
the hydrogen from a sample of each of the successive
electrolytic stages with atmospheric oxygen by
passing the dry hydrogen mixed with dry air over
hot copper oxide at 300°C,
In their second exper­
iment they combined the hydrogen of the successive
stages with electrolytic oxygen from a limited
amount of water and corrected for the decrease in
the density due to the fractionation of the oxygen
isotopes.
From these experiments they estimated
the density of deuterium-free water to be -18,9 Y
and the h /d ratio to be 5600, which is in fairly
good agreement with the value given by Johnston,
Hall and Jones'(9 )
7 corrected for the
fractionation of the oxygen isotopes in two series
of experiments.
First, they equilibrated the water
formed by burning the electrolytic gases of the suc­
cessive stages of electrolysis with carbon dioxide
8,
9,
Morita and Titani,Bull. Chem, Soc.,Japan 11,
403 (1936)
Hall and Jones, J, Am. Chem. Soc.,58, 1915(1936)
(5)
which, in turn, had been equilibrated by passing
it through a large amount of ordinary water.
In
their second experiment the electrolytic gases were
separated and the hydrogen burned in a limited
amount of air.
The density of a sample of the
water from the successive stages was determined, the
water equilibrated with carbon dioxide as in the
first experiment and the density redetermined.
The
density of the water before equilibration in the
separate runs reported differed by 3 p.p.m.
From
the densities of the different waters equilibrated
with carbon dioxide,they calculated the density of
deuterium-free water to be -16.5Y, corresponding
to a h /D ratio of 6500.
The maximum difference between the
densities of deuterium-free water reported in the
literature, even where corrections have been made
for the fractionation of the oxygen isotopes, amounts
to a® much as 2.4Y • This is much greater -than can
be attributed ito experimental errors in density
measurements, and must be due to different methods
used by various authors to bring the water under
investigation to normal isotopic composition.
fe)
Statement of the Problem
Since the relative atomic weight of
oxygen in air and in water is accurately known, it
is possible to solve the oxygen isotope fraction­
ation problem by combining deuterium-free hydrogen
with atmospheric oxygen under the proper conditions,
measuring the density of the resulting water aftd
applying the correction for the difference in the
atomic weight of air and water oxygen*
This would
give the correct density of water composed of pure
protium and oxygen of normal water from which the
correct value of the Tl/D ratio can be calculated*
Experimental Methods
!•
Preparation of Deuterium-free Water*
Seventy-two liters of laboratory
distilled Lake Michigan water, containing four
percent sodium hydroxide in the first and second
fractionations and an equivalent amount of sodium
peroxide in the third, fourth and fifth fractionations,
were successively electrolysed between nickel
electrodes in six cylindrical cells placed in series.
In the first and second fractionations 240 cc, ot
(?)
water were added to each cell and electrolyzed for
twenty-four hours when the residue
and new water added.
was removed
In the third, fourth and fifth
fractionations only 160 cc. of water were added to
each cell and electrolysed twenty-four hours before
removing the residue ahd adding new water.
A current
of 9.5 - 10.5 amperes was used throughout the
experiment.
When sPdium peroxide was used as the
electrolyte, the excess oxygen was removed by boil­
ing.
The electrolytic gases were passed through an
explosion trap consisting of fine sand and were
burned from a pyrex ^et after being dried by means
of concentrated sulfuric acid and anhydrous calcium
chloride.
Bata for the preparation of deuterium-free
water are given in Table I., where the letters
after the numbers are labels serving to distinguish
and identify the various waters mentioned in this
thesis.
( N signifies ordinary , laboratory dis­
tilled, Lake Michigan water, which was chosen as
the standard, and the hydrogen
and oxygen in this
water are assumed to be normal in their isotopic
composition*)
(7a)
Table I
Data Obtained in the Preparation of Deuterium-Free Water
Fractionat ion
Water Water
%
taken, collected collectsd
cc.
<
cc.
<
Residue,
cc.
47,000
cc. retained
Lost, for hydrogen
cc. analysi
1400
X."
72,000(F) 23,600(A) 32.79
2.*
23,600(A)
9,440(B) 40.00
3.
8,640(B)
4,579(C) 52.99
3,782(D)
739
590(C)
4.
3,989(C)
1,865(E) 46.50
1,495(F)
629
565(E)
5.
1,300(E)
690(F) 53.07
530(G)
80
690(F)
800(B)
Table II
Data Obtained in the Preparation of Water Made from
Deuterium-Free Hydrogen and Atmospheric Oxygen
Fraction­
ation
Water
taken
Product,
cc.
Residue,
cc.
Water lost
cc.
200(1)
300
2.
800(B)
300(H)
3.
590(C)
285 (J)
50
265
4.
500(E)
255(H)
50
195
5.
675(F)
440(L)
50
185
*
Mr, Stanley Cristol working under an NYA grant
carried out the first and second fractionations.
2.
Preparation of Water Composed of Deuterium-free
Hydrogen and Atmospheric Oxygen
Part of the water obtained from each
of the fractionations rafter the second, containing
four percent sodium peroxide, was decomposed
electrolytically in a series of three cells specially
designed for the separation of the hydrogen.
The
hydrogen was passed over hot copper to remove any
oxygen present, through calcium chloride to remove
the last traces of moisture,and was combined over
a hot copper catalyst with atmospheric oxygen,
dried by means of concentrated sulfuric acid. Great
care was taken to have the hydrogen always in slight
excess in order to prevent the fractionation of the
isotopes of the oxygen.
The resulting water vapor
was condensed by the use of a water cooled condenser.
In one experiment the condenser cooled by water was
supplemented by pother using dry ice-acetone in order
to determine if there was a detectable fractionation
due to the water cooled condenser, but the cooling
by the dry ice-acetone mixture did not produce a
detectable difference in the density of the product.
Data for the preparation of water consisting of
(9)
deuterium-free hydrogen and atmospheric oxygen are
given in Table II*
3.
Preparation of Rinsing Water
In order to prevent the light water
from being contaminated with normal water, some wash
water for use in rinsing the various glass receptacles,
thermometers, floats,etc., was prepared by electro­
lysing approximately half of the residue from the
second fraction in the hydrogen separation cells
and combining the liberated hydrogen with atmospheric
oxygen.
Two liters of wash water having a density
only slightly greater than that of the lightest
fractions ( J,K,and L ) were obtained.
4.
The Purification and Density Measurements of
the Water
The density measurements were made by
use of the totally immersed float method similar
to that first used by Dufour
lit*
(11)
and improved by
Dufour, Comptes Rendus, 24, 1090 (1862)
Cio)
Richards and Shipleyl^^
The bath consisted of a
battery jar of approximately one gallon capacity
serving as the inner bath, and a larger jar of approx­
imately five gallon capacity insulated with asbestos
serving as the outer bath, in which was placed a
motor driven stirrer.
The density measurements
were made at approximately 30° C.since the temper­
ature of the bath was more easily controlled if the
temperature ofthe laboratory was a few degrees below
the operating temperature.
The bath was heated by
means of a twenty-five watt electric light bulb
controlled by hand, and cooled automatically when
the heater was not in use.
If more rapid cooling
was desired a small amount of tap water was allowed
to circulate through a coil dipping in^o the outer*
bath.
A Beckmanm thermometer was immersed directly
in the water in the measuring vessel, which was a
pyrex tube If-11 by 12w, with the 10 cc. quartz float.
The procedure followed in the actual purification
of the water and density measurements was as follows:
A 250 cc. sample of standard water was completely
11.
Richards and Shipley,.!. Am.Chem. Soc.,34,
599 (1912); ibid, 36, 1 (1914)
(11)
distilled from alkaline permanganate, the steam
condensed after passing over copper oxide heated
to 500 - 600° and the resulting water distilled
from a specially designed flask containing one or
two drops of c.p. phosphoric acid and a few crystals
of potassium permanganate.
The first 50 cc. were
used to rinse the receiving vessel and the next 100
cc. portion for the density measurement.
After the
measuring vessel, Beckmann thermometer, stirring rod,
and -quartz float were rinsed three times with small
portions of the water, 60 - 70 cc. of the water in
the measuring vessel with the quartz float was boiled
vigorously by reducing^ the pressure with a wateraspirator in order to remove completely all the
atmospheric gases from the sample*
The measuring
vessel containing the degassed water, quartz float,
Beckmann thermometer, and the stirrinc rod was then
supported in the inner bath by means of an iron
clamp, and the temperature of the bath adjusted
until th® density of the water was approximately
that of the quartz float.
Ho effort was made to find
the exact equilibrium temperature since such a proced­
ure would have been very tedious.
However, the
(IS)
equilibrium temperature was found in a manner
(IS)
similar to that used by Emeleus and eo—workers,
which consisted of observing the distance of rise
or fall of the float in unit time at several
temperatures on either side of the equilibrium
temperature, and plotting these values against the
temperature*
Over the short temperature ranges
observed, this rate of rise or fall of the float was
found to be a linear function of the temperature,
and the temperature where the curve crossed the axis
was taken as the equilibrium temperature*
In some
of the first experiments, the entire 250 cc* of the
sample were redistilled from alkaline permanganate
and the equilibrium temperature redetermined*
Since
there was no detectable difference in the equilibrium
temperature, the third distillation was considered
unnecessary in the later experiments.
Having thus
established the equilibrium temperature of the float
in normal water, a 250 cc. sample of unknown water
was treated in exactly the same manner as the normal
12*
Emeleus, Jones, King, Pearson, Purcell, and
Briscoe,
J* . Chem. Soc., 1207, 19^8 (193*)
(13)
water, and its equilibrium' temperature determined.
The equilibrium temperature of the normal water
was again determined at this point.
The entire
procedure was repeated at least once for each exper­
iment reported.
Immediately before reading the
thermometer the water was thoroughly stirred, and
the qu^artz float was brought to rest near the center
of the vessel by means of
a tantalum wire loop on
the end of the pyrex stirring rod.
After allowing
time for the float to reach a constant velocity,
the motion was observed for two minutes by the use
of a stop clock and a scale graduated in mm, placed
upright beside the observation window in the bath.
The water was immediately stirred and the final
reading of the thermometer observed.
Since the
average of the two thermometer readings was taken
as the temperature of the water during that
particular observation, the readings were discarded
if they differed more than 0.002° C.
The walls of the float were relatively
thick ( approximately 1 mm. ) and consequently errois
due to different depths of immersion and slight
(14)
atmospheric' changes were not detectable by the
Beckmann thermometer*
The error due to atmospheric
changes was also avoided by determining the equi­
librium temperature in the normal water before and
after each series of experiments.
This is in agree­
ment with the observations of Emeleus and co-workers
However, recent work in this laboratory shows that
slight changes in pressure do affect the equilibrium
temperature when the temperature is being measured
to 0.0001° c.
All glassware, including the float
used in the density measurements, were well cleaned
with cleaning solution and steamed before use, and
were allowed to stand filled,,or in distilled water
between experiments,
^ e n changing from the
standard water to the light waters, the float and
all other parts of the apparatus that came directly
in contact with the water were rinsed thoroughly
with the special rinsing water before starting the
experiment#
An illustration of the data as taken,
is given in Table III.
The quartz float used in these exper­
iments was made and put into use immediately.
Its
(IS)
(14a)
Table III
Ah Illustration of Bata as Taken
Thermometer Readings
Initial
Final
Kotlon of
Float in mm.
Standard Water
3.653
3.651
3.649
3.648
3.648
3.646
3.644
3.645
3.650
3.650
3.650
3.649
3.657
3.658
3.658
3.6585
Equilibrium
5th. Fraction
3.611
3.613
3.685
3.623
3.6195
3.620
3.620
3.620
3.626
3.627
Equilibrium
4 th. Fraction
3.623
3.616
3.624
3.623
3.621
3.619
3.621
3.621
Equilibrium
Standard Water
3.662
3.660
3 *656
3.658
3.650
3.654
3.651
3.650
3.650
3.650
3.658
3.660
3 •659
3.658
Equilibrium
Time
Time of
Bay
-2
3
6
9
1
2
-5.5
-7.5
Temperature
2 min.
2
2
2
2
2
2
2
3.651
10:15
12
-1.5
4
3
4
Temperature
2
2
2
2
2
3.623
11:25
-2
-2.5
1
0
Temperature
2
2
2
2
3.621
12:13
-7
-1.5
1
3
3.5
-6
-5
Temperature
2
2
2
2
2
2
2
3.653
1:55
Aug. 6, 1956
Barometer Reading 751.1 (Corrected).
Beckmann thermometer 3.854 = 30.000° C.
10:25
10:50
11:45
12:30
3:00
0
(15)
©'quilibrium temperature in the normal water increased
rapidly for the first two weeks as shown in Table IT*
(13)
Such observations are reported by Dole,
who
reports an increase in the equilibrium temperature
of 0*047° for a pyrex float over a much longer period,
(14)
and by Richards and Harris,' 1 who report a decrease
in the equilibrium temperature of 0*328°
in
the
first 1752 hours for one of their floats*
5.
Preparation of Light Water from the mixture of
Light Waters
After the density of the various light
fractions had been measured, waters H, J, K, and L
were mixed, the density of the mi±ture was determined,
and a very careful electrolytic fractionation carried
out by decomposing 60$ of the mixture*
In this final
step the air supplying the atmospheric oxygen was
passed through a drying tower of solid
potassium
hydroxide and then through concentrated sulfuric
acid.
The water resulting from the combustion of
the electrolytic hydrogen and the atmospheric oxygen
13*
14*
Dole, J* Am. Chem. Soc*, 58, 580 (1936)
Richards and Harris, ibid, 38, 1000 (1916)
(15a)
Xafcl®
ft*® Xoqwoimm of
%i** Float in
Vi
t $ m
BftsiiXbritig*
T m p ® r & % w m
of
m m terd Water with TAa*
Mt*
BaraNtw
Ju^sr 8* 8*M8
84 m m
747.7
Bat® TbovwMittar Barona)
Hoadis
BoodXsag:
748
A8g* 4 8.0477
747.7
4
3.068
761.7
746.7
a
3.068
761
as 8.a§0
744.3
3.060
761
i? 2.000
74t.f'
6
a
£.863
761
m
m
m
m
*•00*
74a.?
?
8.866
8*910
743.7
750.4
10
IX
3.869
763.4
746.3
3.073
748
86 s**7»
Float
Hftoloaiio*!
jHt
a*9M
a#ft©
3 2.949
4 3.84?
760.4
mu a
?4S
•• BtoXo&socl «* Stood im hot eloaxtlog mlution for
3 hr®. nM stood is the olMtrlo furnace at
aafi* for 18 too**
(16)
was completely recovered in a dry ice-acetone trap.
However, these extra precautions failed to change
the density of the resulting water.
water G
The density of
was measured in order to see whether the
electrodes of the cells had fractionated the oxygen
isotopes during all the electrolyses.
This water
is probably equivalent to Johnstofi!s light fraction
from his second or third electrolysis since it
represents the residue after decomposing 53^ of
water E
from the fourth fractionation.
The ability
of the electrodes of the cells to fractionate the
hydrogen isotopes: was tested at the conclusion of
the experiments.
Twelve hundred cc. of normal
distilled water were placed in the cells and
electrolyzer, and 490 cc. of first fraction resulted.
This sample, water M, had a Y
value of -12.6
which proved that the electrodes had not
,
lost
their activity for hydrogen isotope separation.
Discussion of Data and Comparison
with Previous Results
A temperature of 3.654° on the
Beelmiann thermometer corresponds to 29.8° C.
The
(17)
decrease in density of ordinary water between 29*7°
and 29*8° C. as given by the International Critical
Tables is 50 p. p.m. and can be considered as a linear
function of the temperature*
Thus the
change in
temperature (Zi T) between the equilibrium temperature
of the quarts float in light water and its equilibrium
temperature in norm©.! water offers a convenient
method of obtaining the decrease in density of the
respective light waters from that of the normal water.
The density data are collected in Table V*
The close agreement between the densities
of waters J, K, and L, and the water resulting from
the electrolytic hydrogen from the mixture of these
waters and atmospheric oxygen indicates; that the
hydrogen must have been brought to a constant
isotopic composition which can be assumed to be pure
protium, and that the density of water composed of
pure protium and atmospheric oxygen is 9 p.p.m.
lighter than purified Lake Michigan water.
In order to calculate from -9.oy ,
the density of deuterium-free water containing
normal oxygen, it is necessary to know the difference
in density of waters prepared from both normal
(17a)
Table V
Data for the Density of the Various Waters
Fraction
no*
3
4
5
Label
Ur) -0.029
-
8.7
- .029
-8.7
(K) - .029
-8.7
- .032
-9.6
(L) - .030
-9.0
- ,030
- .030
-9.0
-9.0
- .029
-8.7
Accepted value:
2, 3, 4, 5{Waters B,
Y
AT,°c
K, L> mixed):
2, 3* 4, 5 Mixed and further fractionated (0):
- .030
-9.0
Water G
- .002
-1806
Water M
- .041
-
12.6
(IB)
oxygen and atmospheric oxygen.
This value has
already been determined by a number of worker® using
an electrolytic method with concordant results,
considering the different techniques involved and
the different geographical locations of the standard
waters*
In Table VI*
is a summary of the recently
published results.
By accepting the average value 6.4 Y
as the excess density of water prepared from
atmospheric
oxygen and subtracting this value from
-9.0 Y , we obtain -15.4 Y as the difference between
the density of deuterium-free water and normal water,
the oxygen of which is isotopically identical.
The relation between the decrease in
density of the light waters and the Tl/D ratio in
(8)
the original, or normal water, is
where £* & is the decrease in density,
S the specific
gravity of heavy water, and X is the mole fraction
of heavy wat9r present in the origin?! sample.
The" relation between X
16.
a n d -AS
Lewis andLuten., J. Am. Chem. Soc. ,55,5061 (1955)
Luten, Fhys. Rev., 45, 161 (1934)
f!8a)
Table VI
Difference in Density of Water Containing Normal Oxygen
and Atmospheric Oxygen. (The Hydrogen Having the Same
Isotopic Composition in Each Case)
, [Investigator
Dole(507
6
Greene and Voskuyl (31)
Hall and Johnston (32)
Morita and Titani (33)
Average
Y
6.0
6.0
5 #g
7.0
6.4
Table VII
Density Differences Between Deuterium-Free Water and
ITormal Water from Different Investigations
Investigation
Source of water
Lake Michigan
This investi­
gation
Christiansen, Crab- Melbourne rain
water
tree and Laby
Oxford, England
Edwards, Bell and
Wolfenden
Hall and Jones
Lake Mendota
Ingold, Ingold, Whitaker
and Whytlaw-Gray
London, England
Johnston
Columbus, Ohio
Lewis and Macdonald Berkeley, Calif.
Morita and Titani Osaka, Japan
Tronstad, Nordhagen
and Brun
Bjukan, Norway
r
■15.4
D/H Ratio
1:6900
-12.7
1:8400
1:6200
-16.5
1:6500
-12.0
-18.3
-18.9
1:8900
1:5800
1:6500
1:5600
-18.5
1:5800
(19)
use of the more accurate value, 1,1079, for the
specific gravity of heavy w a t e r , w a s found to be
2.
X Z 9.377AS
- 1.01
If A S in this equation Is the difference
between
the density of heavy water and ordinary water, (Sq ),
the equation can be expressed thus?
5.
X » 9.37TS- 9.377SQ - 1.01S2 +
2.02S§0- 1.01S20
By differentiating the equation with respect to the
partial of S we get: ,
4.
^hx--o~ 9'*377 " 2.02S + 2.02S0
When X reaches its value in normal water the last
two terms become equal antf drop out leavin.fr:
5.
l ^ s ^ o - 9.377
or
By substituting this value and
0,10664
15,4 Y for
A S
■
i
equation 1, we get 6925 for the ratio in normal
Lake Michigan water,
A slightly different value,
7031, would be obtained for this ratio if the
more recent empirical relation recommended by
Longworth^19^
X = 9'.235
17,
18,
l - 0.0327 A S
Selwood, Taylor, Hippie and Bleakney, J. Am,
Chern, Soc,, 57, 642 (1935)
Longworth, J, Am, Chem, Soc,, 59, 1483 (1937)
in
(20)
was rased#
Ira Table Til# is a summary of all the
published results for the Y
values of deuterium-
free water and the h/d ratios rounded off to the
nearest
100
calculated from the
Y
values as above
outlined*
As previously stated it is believed
that the wide discrepancy between these values is
due to the failure of many of the workers to make
corrections for the oxygen isotope fractionation and
to the different methods used in making the correction*
Christiensen, Crabtree and Laby;^-^ Ingold, Ingold,
Whitaker and Whyt law-Gray; ^^^and Tronstad, Nordhagen
(21)
and Brun
in their short M Letters to the Editor*
do not state whether they corrected for the electrolytic
fractionation of the oxygen isotopes in their
(22 )
experiments# Edwards, Bell and Wolfenden
and
(3)
*
Lewis and Macdonald
obtained the H,A) ratio by
determining the rate at which the deuterium
19.
20#
21.
22.
Christiensen, Crabtree and Laby, Nature, 135,
870 (1935)
In gold, Ingold, Whitaker, and Whytlaw-Gray, ibid
134, 661 (1934)
Tronstad,Nordhagen and Brun, ibid 136, 515 (1935)
Edwards,Bell and Wolfenden, ibid 135, 793 (1935)
fsi)
concentrated in cells of known separation efficiencies*
It is difficult to understand how this method could
he as accurate as the density method since there are
so many factors which could change the separation
efficiency*
However, Lewis and Macdonald checked
their values by the density method,
Johnston
' assumed in his work that
the oxygen isotope fractionation is independent of
the kind of hydrogen isotope present in water, but
according to the calculations of Selwood, Taylor,
(17)
Hippie and Bleakney
using the theory of Eyring,
and Sherman^ ^ ^ the Separation ratio for the bonds
IB
16
J> - 0
and D - 0
is slightly higher than that for
18
16
the bonds H - 0
and H - 0 • However, they find
that even in concentratidns of deuterium of 90% the
effect is too small to be detectable
experimentally.
Therefore, unless there are certain unknown factors
Influencing the electrolytic separation, the assump­
tion of the independence of the oxygen isotope
fractionation seems justified.
Probably the
greatest source of uncertainty in Johnston1s work
25.
Eyring and Sherman, J. Chem. Fhys.,1, 345 (1933)
(22)
lies in the extrapolated value of Y
free water.
for deuterium-
A reextrapolation of his data, purposely
sloping the straight line so as to give a value as
near -15.4
, the value of this investigation, as
possible results in the extrapolation value of -15.9 Y*
which is not inconsistent with the results of this
investigation.
However, it is not insisted that this
extrapolation of the data is more reliable than his.
It is difficult to see how any experi­
mental value can be given to the experiment of Morita
(8 )
and Titani
where dry air was mixed with the
electrolytic gases before they were burned, over the
copper oxide catalyst.
It seems that this addition
of atmospheric oxygen would render the results
uninterprstable.
However, the second experiment
where the oxygen from a limited amount of water was
added to the electrolytic hydrogen and burned over
the copper oxide catalyst should have been reliable
and it is difficult to explain the discrepancy between
their results and those' of this investigation.
It
is possible that Japanese water contains more
deuterium than lake Michigan water, but it is unlikely
that the difference is great enough to cause a
difference of 3 p. p.m. in the density of the waters.
Although Hall and Jones,
in their
second series of experiments, obtained results which
vary as much as 3 p.p.m. when they burned their
electrolytic hydrogen with atmospheric oxygen, their
last measurements,
(-8.9 Y)» which were more consist­
ent than the first, agree almost exactly with the
results, ( -9.0
of this investigation*
However,
on equilibration with carbon dioxide this water,
which was apparently exactly like the water, J, K,
L, and 0 of this investigation isotopically, changed
in density to -16.5 y • Whereas, the subtraction of
-6.4 Y 9 the excess density due to the atmospheric
oxygen, would have given
-15.4 y
•
Thus the
discrepancy between their work and this investigation
seems to be entirely due to the two different methods
of correcting for the oxygen isotope ratio.
Their
method of restoring the oxygen isotop© i*atio to its
normal value made use of the isotopic exchange
reactions discovered by lewis
(24)
and was based
upon the following isotopic reaction:
24 '. Lewis, J*. Am. Chem. Soc., 55, 3503 (1933)
Recently, howler, Frey and Greiff
(25)
have shown that in reactions of this type there is
18
an enrichment of 0
in the gas at the expense of the
water, the enrichment factor being 1.039 at 25° C.
for liquid water and carbon dioxide.
They even go
so far as to suggest a method whereby this reaction
1Q
©ould be used as a means of concentrating the 0
in water.
Applebey and O g d e n ^ ^ demonstrated this
principle by taking a sample of purified normal water
and bubbling a large excess of sulfur dioxide through
it.
After
repurification and distillation, its
density was 3 p.p.m. lighter than the original water.
These results agreed well with the calculated value
of 2.3 p.p.m. from the enrichment
factor, 1.014,
at 25° C. for the reaction
S0216+
H2018
S0218
HgO16
Therefore, it is believed that the method employed
in this investigation,which consists essentially
in measuring the oxygen isotope ratio in atmospheric
oxygen, gives a more reliable result than the carbon
25.
26.
Frey and Greiff, «T. }Am;:Chem. Soc., 57,_321(1935)
Applebey and 0«?den,
GhemI* B o h 1^-1835 (1936)
dioxide equilibration method used by Hall and Jones
Calculations of Some
Equilibrium
Isotopie Exchange
Constants *
In the work of Hall and Jones the
assumption is also made that the oxygen isotope
equilibrium between carbon dioxide and water is the
same irrespective of the deuterium content of the
water.
Therefore, it will be of interest to calculate
the equilibrium constants of the oxygen isotope
reactions between carbon dioxide, oxygen, and heavy
water to see if this assumption is valid.
When we know the distribution function
ratio,
the equilibrium constants
for the reactions
16
18
.
18
16
6. C02 -h 2Dp0 ( g) ^ C 0 g H- 2Dg0
(ff)
7.
COs16-f- 2D2018 (1)
S.
02
9»
0216+ 2D2018 (|J
#
16
+• 2 D 20
18 . | /
(g) ± = 9
COg18 -f 2D2016 (1)
02
18 ,n, _16 . .
-h
2DgO
(g)
0818 -h 2D2016 C>)
The original calculations were made by Dr. H.Dole/
Chemistry Dept., Northwestern University,Evanston,111.
(26)
Can be calculated by makidg use of other distribution
ratios tabulated by Frey and GTeiff^^
10-.
U
u '.
K s
-
6-f C O ? / f \ o “ f C O ?
^foL’f i L o ' y / c ?
/
a
o
'«
The distribution function ratio for poly-atomic
molecules is given by the following equation;
Where A, B, and C are the moments of
inertia of the molecule with respect to the principle
axis through the center of mass, uSl is the fundamental
frequencies of the molecule , J I
means that the
product is taken over all frequencies of vibration
of the molecule, c3i , one for each vibrational degree
of freedom, and the other constants have their usual
meaning#
The frequencies of the BgO16 and T>gO*^ were
calculated by the equations:
13,
14 .
U.TT^a^
(*„ ~
«'0
J> + A " ' ^ ' °
(27)
15. £(Crh
'ki -Ati-J Jjiii )£-
-f-
16.a„= «lu= H _ !a
17.
-2.[q<i/V*~/f'3_J- O
__ _
^ ^+
I
Z0 V ^ 3
_
„
/T-
b
^
e
18. <*33 - A ^ Z ^ J )
.
^
/f»^vv ^
19. 3,} - Z5— — — -—
•2. ire
--
used by Van Vleck and Gross
frequencies of the
(27)
in calculating the
molecule based on quantum
mechanics and spectroscopic- data from the OH and Hg
molecules#
They evaluated the force constants, k«^,
kig, kg, and k^g, by use of the potential energy <5>^
the HgO molecule according to the quantum theory of
directed valence, and the a.*£ constants by equations
(16) to (19) given above derived from the t*agrangian
function of the HOH molecule where
are the reduced masses and &
27.
Ch-h^) and A 3
is the apex angle
Van Vieck and Cross, J. Chem. Phys., 1, 350 and
357 (1933)
of the molecule*
Since the BOD molecule may be assumed
to have the same potential function as the HOK .
molecule, the fundamental frequencies
and 2784 of the molecule
2666* 1179,
observed by Rank, Larsen,
and Bordner^8 ^ and Barker and Sleater^^ may be
used in equation (13) to evaluate the constants XX*
X X $ and
values of the
• These values and the calculated
constants used in equations (14)
and (15) give three equations which solved simul­
taneously will give the new values of the force
constants k^, k-^g, and K3 provided the value of k^g*
which is small compared with the others, ( In the
HOH molecule k]_g s0#04 x 10^) is given the value,
0*05 % 10 #
With the force constants for the
BOD molecule the desired frequencies can be
calculated, and the distribution function ratio
determined#
The results of the calculations are
shown in Table VIII#
The distribution function ratio,
28.
F*»nk, Larsen, and Bordner, J. Chem. Fhys.,2,
464
(1934)
29. Barker and Sleater, J. Chem, Phvs.,3, 660 (1935)
(28a)
Table VIII
Force Constants of DgQ Molecule
^1
^3
^12
7.796 X ID5
0.4976 X 105
^13
0.3698 X 105
0.05 K IQ5
Molecular Constants
Molecule
D2016
IB
D2Q
Wg
Ag/Ai
2666.0
2785.0
1176.6
2646.8
2770.3
1169.5
Bg/Bi
1.0083
1.0229
Distribution Function Ratios
fDsO18
fD20-Lti
273,1°A.
1.3043
298.1°A.
1.2926
600°A.
1.2316
Equilibrium Constants and Enrichment Factors
Reaction
16
C02
4“ 2D201£<.S)
18
COs-^ -H 2D2016(g)
16
C02X
t EDgO^d)
16, .
2DgO (1)
CO^
o216
A
~h
sxtfpf-jg)
18
2D2016(g)
A
021 6
2D2018C1)
o
°218
2D2016(1)
02
~h
&298.I
r
Enrichment Factors
1.065
1.032
1.036
1.018
0.9990
00.9995
0.9716
.9857
University
LM>rar»
f A O * / f o * 0 "ratio *j~
f is about 2io greater than the
fz-tiO*** , hence the equilibrium
constants of the reactions listed in Table VT. are
about 4 $ smaller than the constants for the normal
18
water equilibria
Thus 0
tends to concentrate
less in the gas phase and more in the aqueous phase.
In the case of the oxygen-water equilibrium the new
calculations show that the equilibrium concentrations
1ft
of the reactions are shifted far enough so that 0
concentrates in the heavy water rather than in the
gas*
However, the difference in behavior between
the carbon dioxide-heavy water equilibrium is not
great enough to invalidate the equilibrium method
of Hall and Jones.
In regard to the relative rates
at which the carbon dioxide comes into equilibrium
with the two kinds of water, nothing is known at the
present time.
Summary
Deuterium-free hydrogen has been prepared,
and combined with atmospheric oxygen yielding water
lighter by 9 p.p.m. than Lake Michigan water.
6.4
When
p.p.m. are added to 9 p.p.m. to correct for the
difference in atomic weight of atmospheric oxygen
and aqueous oxygen, the density of deuterium-free
water containing normal oxygen is
lighter than normal water.
15.4 p.p.m.
Prom this value the
ratio of deuterium atoms to hydrogen atoms in Lake
Michigan water is calculated to be 1 : 6900, and
confirms the belief of Hall and Jones that the
commonly accepted value for the T>/E ratio is too
high and should be revised downward.
The result
indicates, however, that the downward revision
should be somewhat more than that recommended by
Hall and Jones.
Isotopic exchange equilibrium and
16
18
force constants involving Ds0
and DgO
have
been calculated and tabulated.
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I
A.
Isotopic Studies
A Redetermination of the Protium-
deuterium Ratio in Normal Water
Urey, Harold C., F.G .Brickweddef
G*M>’::>Murphy
A Hydrogen Isotope of Mass 2.
Physical Review, 39, 164 (1932).
A Hydrogen Isotope of Mass 2 and Its Concentration
Physical Review, 40, 1 (1932).
Birge, Raymond T. and D.H.Menzel
The Relative Abundance of the Oxygen Isotopes, and
the Basis of the Atomic Weight System.
Physical Review, 37, 1669 (1931).
Bleakney, Walker
A Search for Isotopes of Hydrogen and Helium.
Physical Review, 41, 32 (1932).
Hardy, J.D., E.F.Barker and D.M.Dennison
The Infrared Spectrum of H^Cl.
Physical Review, 42, 279 (1932).
Tate, John T., and Philip T. Smith
An Attempt to Observe a Helium Isotope.
Physical Review, 43, 672 (1933).
(32)
-5.
Bleakney, Walker and Austin J.Gould
The Relative Abundance of Hydrogen Isotopes.
Physical Review, 44, 265 (1933).
6.
Lewis, Gilbert N. and Ronald T.Macdonald
2
Concentration of Hq Isotope.
The Journal of Chemical Physics, 1, 341 (1933).
7.
Johnston, Herrick L.
The Preparation of Deuterium-Free Water.
Deuterium
Content of Ordinary Water and the Atomic Weight of
Hydrogen.
Electrolytic Separation of the Oxygen
Isotopes.
The Journal of the American Chemical Society
57, 484 (1935).
8.
Morita, Noriyoshi and Toshizo Titani
Erzeugupg des Leichten Wassers und Bestiramung der
Deuterium-Konzentration im Normalen Wasser.
Bulletin of the Chemical Society of Japan
11, 403 (1936).
9.
Hall, Norris F. and Thomas 0.Jones
A Redetermination of the Protium-Deuterium Ratio
In Water.
Journal of the American Chemical Society
58, 1915 (1936).
(33)
10.
Dufour, Comptes Hendus.
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11.
Richards, Theodore W. and John W. Shipley
A New Method for the Quantitative Analysis of
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Journal of the American Chemical Society
36, 1 (1914).
12.
Emeleus, H.J., F.W.James, A.King, T.G.Pearson
R.H. Purcell and H.V.A.Briscoe
The Isotopic Ratio in Hydrogens
A General Survey
by Precise Density Comparisons Upon Water From
Various Sources.
Journal of the Chemical Society, 1207 (1934).
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13.
Dole, Malcolm
The Concentration of Deuterium In Organic Compounds.
II. A General Discussion With Particular Reference to
Benzene.
Journal of the American Chemical Society
58, 580 (1936).
(34J
14.
Richards, Theodore W. and Gorham W. Harris
Further Study of Floating Equilibrium,
journal of the American Chemical Society
38, 1000 (1916).
15.
Lewis, Gilbert N. and Daniel B. Luten, Jr.
2 18
The Refractive Index of HoO, and the Complete
Isotopic Analysis of Water.
Journal of the American Chemical Society
55, 5061 (1933).
16.
Luten, Daniel B. Jr.
The Refractive Index of
and Density of Solutions of
The Refractive Index
2 2
1 1
Physical Review, 45, X6X (1934).
17.
Selwood, P.W., Hugh S. Taylor, J.A.Hipple, Jr. and
Walker Bleakney.
Electrolytic Concentration of Oxygen Isotopes.
Journal of the American Chemical Society
57, 642 (1935).
18.
Longworth, L.G.
The Densities of Mixtures of Light and Heavy Water,
Journal of the American Chemical Society
59, 1483 (1937).
(35)
19.
Christiensen, Ur.N., R.W. Crabtree, aSdHT.l$Sfo;pLaby
Density of Light Water: Ratio of Deuterium to
Hydrogen in Rain Water,
Nature, 135, 870 (1935).
20.
Ingold, E.H., C.K. Ingold, H. Whitaker, aftd
LIew-
Wbytlaw-Gray
Preparation of Protium Oxide and Determination of
the Proportion of Deuterium in the Hydrogen of
Normal Water.
Nature, 134, 661 (1934).
21.
Tronstad, L,J.Nordhagen, ahd .
Jjsv©run
Density of 100 per cent Heavy Water.
Nature, 136, 515 (1935).
22.
Edwards, A.J,? R.P. Bell, ahdl^.BsIWolfeBden
Deuterium Content of Naturally Occurring Water.
Nature, 135, 793 (1935).
23.
Eyring, Henry and Albert Sherman
Theoretical Considerations Concerning the
Separation of Isotopes.
Journal of Chemical Physics, 1, 345 (1933).
(36)
24.
Lewis, Gilbert H.
A Simple Type of Isotopic Reaction.
Journal of the American Chemical Society
55, 3503 (1933).
25*
Urey, Harold C. and Lotti J. Greiff
Isotopic Exchange Equilibria.
Journal of the American Chemical Society
57, 321 (1935).
26.
Applebey, Malcolm P. and Geoffrey Ogden
The Electrolytic Preparation of Deuterium and
the Separation of Coefficient a.
Journal of the Chemical Society, 163 (1936).
27.
Cross, Paul C. and J.H. Van Vleck
Molecular Vibrations of Three Paricle Systems
with Special Applications to the Ethyl Halides and
Ethyl Alcohol.
Journal of Chemical Physics, 1, 350 (1933).
A Calculation of the Vibrational Frequencies
and Other Constants of the HgO Molecule.
Journal of Chemical Physics, 1, 357 (1933).
28.
Rank, D.H., K.D. Larsen and E.R. Bordner
The Raman Spectrum of Heavy Water Vapor.
Journal of Chemical Physics, 2, €64 (XS34).
(37)
29•
Barker, E.F. and^W.W. Sleator
The Infrared Spectrum of Heavy Water,
Journal of Chemical Physics, 3, 660 (1935).
30.
Dole, Malcolm
The Relative Atomic Weight of Oxygen in Water and
in Air a Discussion of the Atmospheric Distribution
of the Oxygen Isotopes and of the Chemical Standards
of Atomic Weight.
Journal of Chemical Physics, 4, 268 (1936).
31.
Greene, Charles H. and Rodger J. Voskuyl
An Explanation of the Relatively Large Coneentra18.
txon of 0 xn the Atmosphere.
Journal of the American Chemical Scoiety
58, 693 (1936).
32.
Hall, V/. Heinlen and Herrick L. Johnston
Influence of Combustion Conditions on the Density
of Water Formed from Commercial Hydrogen and Oxygen.
Journal of the American Chemical Society
58, 1920 (1936).
33.
Morita, Von Horiyoshi and Toshizo Titanx
Erzeugung des Leichten Wassers und Bestimraung der
Deuterium-Konzentration im Hormalen Wasser.
Chemical Society of Japan Bulletin
11, 403 (1936).
(1)
Micro determination of the Density of
Small Quantities of water by the
Balanced Drop Method
Review of Literature
Of the various methods used for the
quantitative analysis of water for its deuterium
content
(l)
the method based upon the density of the
water is the most suitable for laboratory use, (see
above 3T-A, P*18
is low*
) especially when the concentration
The pycnometer method, the standard method
for measuring the density of water with extreme
care, has an accuracy of 1 p.p.m. or an error of
0.001$ in deuterium concentration.
However, the
method is tedious, time consuming, and requires
a relatively large sample of waterr
The totally
submerged float method, based upon the difference
in the density-temperature coefficient of the
float and that of the water, first devised by Dufour
1. Urey and Teal, Rev. of Mod. Phys., 7, 54(1035)
2'. Dufour, Comptes Rendus, 24,1080 (1862)
(2)
and improved by Richards? and co-workers
(3-4;)
is
unsurpassed in accuracy and due to its sensitivity
and precision has come into general use.
The
method as employed in this laboratory, where the
temperature is measured accurately with a platinum
resistance
thermometer to 0.0002s0 C. , has
an
accuracy easily within 0.2 p.p.m., or 0.0002$
deuterium.
in
Many modifications of this method are
use in various laboratories. Lamb and Lee (5)
fastened a permanent magnet to the float and made use
of small platinum weights and a magnetic coil to
balance it.
The current through the coil was reduced
until the voltage was found which
float from sinking*
.fust
Prevented the
Hall and Jones' 7 used this
method to measure the density of the water in their
(7)
experiments.
Gilfillan and Polanyi
made use of
the Cartesian diver based upon the difference in
in the density-pressure coefficient of the water
and the float* and determined the density of the
3.
4.
5.
G.
7.
Richards and Shipley, J. Am. Chem. Soc.,34,599(1912)
ibid
36,1,(1914)
Richards and Harris, ibid
38,1000 (1916)
Lamb and Lee, J. Am. Chem. Soc.,35, 1666(1913)
Hall and Jones, J. Am. Chem. Soc., 58,1915(1936)
Gilfillan and Polanvi,Ziets.fur Physik Chemie
166,254 (1933)
(3)
water from the pressure at which the diver lust stink.
(8)
(9)
@ilfillan
and Greene and Voskuyl
applied the
same principle to closed floats, using a thin walled
float which was more compressible than the water for
large density differences, and a thick walled float
less compressible than water for small differences.
Relatively large samples of water are used in the
usual laboratory procedures for the purification
and density measurements by all these methods, and
in this laboratory a standard sample of 250ec. was used.
However, in the use of deuterium as an indicator in
the study of biological and chemical problems it is
essential to be able to measure the density of very
small quantities of water.
Hittenberg and Sehoenheimer’^
have improved the procedure based on the densitypressure principle with only a small decrease in
accuracy to samples as small as 2.5 cc.
ibility
upon four samples of triply distilled
water was within 1.13 p.p.m.,
and Thomas
The reproduc­
Rromhez,Sonderhoff
have succeeded in reducing the volume
S. Gilfillan, J. Am. Chem. Soc.,56,406 (1934)
9. Greene and Voskuyl, ibid
56,16^9'(1934)
10. Rittenberg and Schoenheimer, J. Biol. Chem.
Ill, 169 (1935)
11. Fromhez,Sonderhoff,and Thomas,Berichte,70,1219(-1937)
(4 )
of the sample to 1 cc. with an accuracy of 1.3 p.p.m.
or 0.0014$ deuterium.
In their experiments the
temperature was varied with the pressure constant
until the float was
approximately balanced and then
the pressure was varied with the temperature constant
until the float was exactly balanced.
This procedure
is more rapid since the pressure can be changed with
much greater rapidity than the temperature of the
entire bath.
The volume,! cc., is undoubtedly about
the minimum necessary for the density measurements
by the float method unless the sample is diluted, and
this is undesirable since the error in diluting is
much greater than that in the density measurement
(12)
itself.
Fenger-Erieksen, Krogh, and TJssing
have
recently improved the ingenious method of Barbour and
)
Hamilton'( 1 3 1
for measuring the specific gravity of
body fluids and report an accuracy of 2 p.p.m. when
the method is used in determining the density of
waters containing various known percentages of deuterium
12.
13.
Fenger-Ericksen, Krogh, and Ussing, Bio.,Chem. J•
30, 1264 (1936)
B a r b o u r , a n d Hamilton,!. Bio. Chem.,69.625(1926)
Am. J. Physiol. EXIX, 654 (1924)
(5)
oxide*
Their procedure was as follows:
In a
constant temperature bath were placed seven glass
stoppered tubes 45 cm* long and 1*6 cm* in diameter
containing bromobenzene and xylene mixed in such
proportions that a carefully measured drop of pure
water and drops of water containing 0*5, 1.0, 1.5,
2.0, 2*5, and 3*0 percent deuterium oxide would fall
a distance of 9 cm. through the respective tubes in
about 30 seconds.
Four grains: of anhydrous sodium
sulfate were placed in the bottom of each tube to
take up the water.
A drop of water of unknown
deuterium contemt delivered by a carefully designed
micropipette was allowed to fall through the mixture
where the rate was approximately 9 cm. in 30 seconds.
The rate of fall of the drop was timed at 9 cm. inter­
vals over a distance of 27 cm..
The average time
for the drop to fall throufh 9 cm.
was compared
with that required for a drop of water of known
deuterium contrent to fall the same distance, and
the deuterium content of the unknown sample calcu­
lated by assuming that the density was a linear
inverse function of the time.
Hochberg and
(6)
LaMer
not a
(14)
however, have shown that the density is
linear inverse function of the time, and that
the actual linear interpolated value of the illus­
tration cited by Eenger-Ericksen and co-workers was
22 p.p.m. greater than their reported value.
These
authors also have shown by using the relation
d C k/t - d'
where d is the density of a drop fallin& through
a fixed distance in time, t, and k is a constant
depending on the viscosity and density, d', of the
medium and the distance over which the dr6p is timed,
\
that it is possible to calculate the density of a
water solution using only two drops of the solution
of 0.001 - 0.01 cc. volume.
The falling time of
the drop is compared with that of two other drops
of known but different densities over a 15 cm.
distance through the bromobenzene-xylene mixture.
By solving the three equations simultaneously it is
possible to obtain the density ,,of nthe , unknown
sample to i 0.00001.
14.
To obtain an accuracy of
Hochberg and LaMer, Ind. and Eng. Chem. Anal.
Ed. 9, 291 (1957)
(V)
0.00001 it is necessary in making the interpolations
to apply corrections, the magnitude of which depends
upon the velocity differences, and hence the density
differences over which the interpolations are made.
Statement of Problem
The large error in the Palling Drop
Method is due to the failure of the density to follow
linearly the inverse function of the time.
It is
reasonable to believe that this is due to the direct
relation of the distortion of the drop to its velocity
through the medium*
Therefore, a method of balancing
the drop of water in an immiscible medium by use
of the difference in the density-temperature
coefficients of the medium and the water by a change
of temperature as in the totally submerged float
method would eliminate this error and should provide
a highly accurate method of measuring the density
of a single drop of water provided the purification
of such a small quantity of water is reproducible.
Some of the properties of an ideal
medium would be:
A density equal to that of water
at some temperature between 0° an(3 30°
a small
(8)
density-temperature coefficient, a low viscosity,
low volatility or high boiling point, complete
insolubility of the water and oil in each other, and
easily ,repurified.
Since slight evaporation would change
the density of a mixture and repur1fixation is not
practicable, a pure compound should be selected if
possible.
Several relatively pure compounds were
obtained with the correct density such
anisole,
n-butyl benzoate, o-tolunitrile, ethyl benz?rl ketone,
1**4-dihydronaphthalene, and cyclohexyl chloride.
Because of their solubility, the compounds contain­
ing oxygen and nitrogen were eliminated*
1
^he
4-dihydronaphthalene proved to be unstable with
the equilibrium temperature of the drop changing
rapidly with time.
Prom preliminary experiments
upon a sample obtained from the Eastman Kodak
Company, cyclohexyl chloride seemed to be a suitable
medium and was selected for this investigation.
It should be mentioned that while this
investigation was under way, Keston, Bittenberg,
(15 )
and Schoenheimer
improved the Palling Drop Method
15. Keston, Rittenberg, and Schoenheimer, I. Bio.
Chem.,122, 227 (1957)
(9)
by
using the middle fraction of o-fluorotoluene
with a density of 0.9995 at 25*6?®’C* for the medium*
By increasing the falling time of the drop of pure
water over a 15 cm, distance to 180 seconds, they
were able to compare two drops with an accuracy of
1 p#p*m*,
The deviation of the density from the
inverse function of the time varied less than
0*01 atom % up to 1*0 % deuterium oxide concentration*
They checked their density measurements with a Zeiss
interferometer calibrated to read directly the
percentage of deuterium oxide present in the sample
according to the method of Crist, Murphey, and Urey*
(10)
They found the method rapid and very suitable for
routine laboratory procedures*
Materials and Experimental Procedure
1* Preparation of Cyclohexyl Chloride
The cyclohexyl chloride was prepared,
with slight modifications, according to the directions
16.
Crist, Murphey, and TTrey, J. Am# Chem# Soc#,
55,5060 fl933)*
J, Chem# Phys., 2, 112 (1934)
(10)
£n Organic Syntheses
n-butyl chloride*
for the preparation of
Five moles of commercial oyclo-
hcxanol was refluxed
ten hours with ten moles of
hydrochloric acid of 1*19 specific gravity and ten
moles of anhydrous zinc chloride*
After cooling,
the oily layer was separated from the water solution
and washed several times with water*
The oil was then
washed with cold concentrated sulfuric acid*
If the
oil was heated with concentrated sulfuric acid,conden­
sation
tools place, or if allowed to stand too long
in contact with cold concentrated acid, the line
between the two phases became invisible*
The oil
was then washed free of acid with watery dried with
calcium chloride, and distilled from a Claisen flask
under reduced pressure*
The second distillation
gave approximately 3*00 cc. of the oil with a constant
boiling temperature.
The oil hoiled at 142° C* under
atmospheric pressure, and decomposed slightly if
heated much above that temperature*
2>* Measurements of the Density-Temperature Coefficient
17.
Organic Syntheses, Collective Tol* 1, 137*
John Wiley and Sons, New York (1932)
(11)
of the Oil Saturated with Standard Water
Dr* F. T'. Gucker Jr. very graciously
permitted the use of his thermostat designed for
cis)
pyenometer work for these density measurements*
The temperature was measured by a platinum resistance
thermometer and held constant to “± 0*0004° C.
The
pycnometer was of the ordinary 10 cc. capacity type
modified slightly for these experiments*
The deliv­
ery end of the pycnometer was drawn out to a fine
capillary, a capillary tube with a single graduation
mark and a small bulb to allow for expansion of the
liquid replaced the other arm of the pycnometer,and
the ends were covered with close fitting caps to
minimise evaporation*
The weighings were made on a semiml6l?o-balance with a notched beam graduated in hun­
dredths of a milligram and magnetically damped. The.
weights had been standardized by Gucker and co-workers
(19)
according to the method of Richards*
An empty
pycnometer of apprpximately the same capacity was used
18*
19.
Gucker.,-, Gage, and Moser, J* Am. Chem. Soc.,
60, 2582 (195 8)
Richards, J. Am. Chem. Soc., 22, 144 (1900)
(12)
as a tare in order to decrease the error due to
adsorption of moisture on the surfacer all weights
were reduced to the vacuum standard, and corrections
were made for the cubical expansion of the pycnometer.
After the thermostat had reached equilibrium at- 20° C.,
the pycnometer with a slight excess of oil saturated
with standard water was placed in the bath.
Twenty
minutes were allowed for the temperature to reach
equilibrium, the meniscus was carefully adjusted
in the usual manner by use of a small piece of filter
paper, and after another ten minutes the final adjust­
ment of the meniscus was made.
The pycnometer was
removed from the water bath, rinsed with distilled
water, dried with a lintless towel and hunc? in the
balance case.
manner.
The tare was treated in the same
After approximately thirty minutes, three
successive weighings were made at fifteen minute
intervals.
& slight excess of oil was then adde'd
to the pycnometer and the procedure repeated.
The
same procedure was followed at the temperatures
20.497° and 20.9943° 0*. The volume of the pycnometer
was found by filling it with the purified standard
water and weighing in a similar manner.
The data for1
(12a)
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(13)
the density-temperature coefficient measurements
are given in Table I*, and the absolute density of
the saturated oil and water are plotted against the
temperature in Fig*la, while the net difference
(
in the density of the two liquids are
plotted against the temperature in Fig. lb.
3.
Some Preliminary Experiments
A thermostat constructed by Mr. R.L.
Slobod was used for the preliminary experiments, and
it consisted of two large circular tubs, one inside
the other.
The inner tub was covered with sheet
copper containing holes for the measuring vessel and
the Beckmann thermometer.
The measuring vessel, a
pyrex tube 50 cm# long and 1 cm. inside diameter,
was placed in the exact center of the thermostat by
means of a rubber stopper fitted into a hole in the
wood cover.
The Beckmann thermometer was placed
just outside the measuring vessel in the inner bath.
The outer bath contained the motor driven stirrer,
•two electric heaters, one for rapid heating and one
for slow heating, and the copper cooling coil through
which tap water could be circulated.
The temperature
(13 ft)
2ao
2 0 .4
Pig. la. Absolute density of the saturated oil
and water as a function of ternnenature.
400
300
CL
/oo
Fi^r. lb. Net difference in theX
densities of the saturated
oil and water.
3oo
2 o.o
(14)
was automatically controlled to dz 0.001° C. by
means of a raercury-toluene regulator operated by
(20)
a vacuum tube relay,
several of which are in use
in this laboratory*
The thermostat was designed to
operate at one or two degrees below room temperature
and contained no insulation around the sides*
The purified hake Michigan water IT,
which served as the standard water, was purified
as described in part one of this thesis* ( see above
I-A,
P. lo)
A pyrex tube 5-6 mm* in diameter drawn
to a fine capillary and fitted with a dropping bulb
served as a pipette for ejecting a small drop of the
purified water Just below the surface of a 15 cc*
sample of the oil*
The water clung to the glass
and was slowly pulled from the glass tip by the sur­
face tension of the oil.
The motion of the drop
was then observed by the use of a cathetometer*
The equilibrium temperature of the drop
of water in the oil was obtained by a procedure
similar to that used in the submerged float method.
«
(see 1-A,p*9.)# The distances which the drop rose
20.
Thiessen and Frost,J. Chem. Edc.12,72 (1935)
or fell in three minutes, observed at several
temperatures on dither side of the equilibrium
temperature, were plotted against the temperature
and the equilibrium temperature found with the aid
of the graph.
This procedure
made the equilibrium
temperature independent of the size of the drop.
An illustration of the data is given in Table XI.
and plotted in Fig. 2#
It was found that a very rapid decrease
of the equilibrium temperature, amounting to a few
hundredths of a degree, followed the introduction of
a drop of the purified standard water into the dry
oil.
Significant data covering this change were
impossible because of the rapidity of the change.
In subsequent experiments oil saturated with standard
water was used.
When a drop of water was added to
the saturated oil slightly below its equilibrium
temperature, the drop would fall slowly for three
or four minutes and then slowly rise with a constant
velocity.
This difficulty could be partly eliminated
by adding a small drop of oil to the water before
the measurements were made.
This was an indication
that the water-oil mixture was slightly lighter than
ClEtf
Table IX'.
Bata Illustrating a Typical Measurement
________
Water G (Y = -6.5)_________________
Tima
Temperature
9:05
9408
9:11
9:14
9:17
9:20
1*938
1*939
1.940
1.941
1.941
1.941
10:09
10:12
10:15
1.930
1.9295
1.9295
10:25
10:28
10:31
1.931
1.931
1.931
30.66
30.76
30.80
30.76
30.80
30.85
11:00
1.933
1.934
1.934
1.935
1.936
1.936
1.936
1.937
31.61
31.58
31.48
31.32
31.22
31.05
30.92
29.79
31.58
31. 48
31.32
31.22
31.05
30.92
29.79
29.63
11:03
11:06
11:09
11:12
11:15
11:18
11:21
j
Cathe tometer* Headings
Initial
Pinal
Distance
cm. in 3 min.
32.85
32.63
32.36
31.98
31.55
31.16
29.97
30.11
30.25
32.63
32 .36
31:98
31.55
31.16
30.73
30.11
30.25
30.39
-
0.22
-0.27
-0.38
-0.43
-0.39
-0.43
*0.14
*0.14
*0.14
* 0.10
*0.04
*0 .05
-
0.02
-
0.10
-0.16
- 0.10
-0.17
-0.13
-0.13
-0.16
(15b)
+ 3 0
<•
1926
/9Z8
/.93 0
/.93Z
193+
/9 3 6
/.ff3Q
'940
/.94Z
/.94+
Fig. 2
Equilibrium temperature of d water in the saturated oil
(16)
either the pure oil or the pure water*
If the drop of water came in contact
witljj: the glass, there was danger of its spreadingon the glass surface, especially if the difference
between the temperature and the equilibrium tempera­
ture was great*
Therefore, great care was exercised
to prevent the drop from coming in contact with the
glass surface*
In spite of these precautions drops
were often lost in this manner*
In later experiments
the drops were removed with the pipette immediately
following the measurements.
In case the drops
were water of abnormal density and were lost, or
when several drops of the N water were lost, the oil
was placed with the wet oil and later dried with
magnesium perchlorate, redistilled and resaturated
with the BT water for future use*
The data covering
the preliminary experiments are given in Table III*
It can be seen from the data in Table
III* that the reproducibility of the equilibrium
temperature
of the standard water in the oil on
successive days was fairly good.
It should be
noted that occasionally there was a relatively large
rice in the equilibrium temperature ‘probably due to
(16a)
Table III.
Equilibrium Temperatures of Water IT and Standard
Light Water g In Oil in the Preliminary Experiments
History cf
the oil
3/24/38 Saturated with
H water
Bate
Water Equilibrium
temperature
H
1.500
3/25/38 Excess water
removed by
centrifuged.
3/26/38 Same oil
H
<?
1.501
1.526
3/30/38 Same oil
<J
1.547
3/30/38 Same oil
H
1.535
Hew oil
saturated with
H water
H
1.501
4/4/38
Same oil
H
1.529
4/6/38
Hew oil
saturated with
H water
H
1.533
Hew oil
saturated with
H water
H
1.529
Last two
samples mixed
H
1.531
Same oilExcess water
removed by
centri fuge:
Divided— -- (a)
<3
1.541
(b)
<?
1.566
4/1/38
4/6/38
4/8/38
4/8/38
See below
p. 18
Remarks
Hote change in
equilibrium T.
A T -~-0. 0/2-
Large change in
equilibrium T.
A t
o .oio
Large changein
equilibrium T.
(vn
some unknown contamination of the oil*.
Subsequently
trhe measuring vessels and all other flasks containing
the oil were fitted with ground glass stoppers and
guAarded carefully against contamination*
Also
considerable difficulty was experienced in controlling
the temperature of the thermostat to within 0*001°
o
at approximately 20.5 C*, especially when the room
temperature was approximately 30a C*j so a new ther­
mostat was constructed for the more accurate density
measurements*
A thermostat of the same general design
as the one previously described was constructed*
The bottom and sides, except for narrow observation
windows on each side,were insulated by a three eighths
inch wood frame plus three inches of magnesia insulation,
and the top by a one fourth inch plywood cover plus
two layers of insulating matting*
Ho*
Two five-ohm
Uichrome wire heaters, operated by the
electric light circuit, reduced to fifteen volts by
a variac of fifteen-ohm capacity, heated the bath*
At first tap water circulating through a copper coil
around the inner bath was used for cooling the ther­
mostat.
However, it was later found that a much
(18)
more uniform cooling could be obtained by dropping
ice water into the bath at a uniform rate,
A motor
mounted on the brick wall above the thermostat
furnished adequate stirring, and all vibrations
were practically eliminated by placing the thermo­
stat upon a solid concrete pier resting on a concrete
floor,
4*. Density Measurements
Two samples of light water prepared, in
this laboratory by R. L. Slobod and J. A, Swart out
served as the standard light waters.
Water GT was
composed of deuterium-free hydrogen and fltrosviXle.
Limestone oxygen, and water B was twice electrolysed
Lake Michigan water* ( Similar to water B given in
Table II* i-A, P* 8a)
All waters were purified
in the same manner as previously described,
above I-A, p, 10)
fsee
The exact densities of the light
waters were obtained by the totally submerged float
method as described in this paper with certain improve­
ments*
($ee above I-A,tP«# 0 ) gh@ measurements were
made on the thermostat just described where the
temperature was accurately controlled
to
0.001° C,
(19)
Since the equilibrium temperature of the quartz
float used was between 24.0°
and 25.0° C. where
the density-temperature coefficient of water is
0.23 p.p.m. per one thousandth decree, the measure­
ments are accurate to ± 0*3 p.p.m.
A summary and
comparison of the density measurements are given
in Tables tV and V.
5*
Solubility of the Oil in Water
Two 300 cc. samples of purified standard
water containing a slight excess of the oil was
allowed'to stand for several days (2-4).
One hundred
cc. samples of the water were removed and refluxed
with an excess of alcoholic silver nitrate for two
hours*
The precipitates were collected
in weighed
Gooch crucibles, dried in the oven at 160° C*, and
weighed as silver chloride.
In pipetting the
samples care had to be exercised to prevent fine
drops of the oil from being sucked into the pipette.
In one case the water was allowed to stand in a tall
pyrex cylinder and the sample pipetted from near the
center.
Two samples checked each other within
and checked
0.4 mg.
the previous determination within 0.9 mg.
(19a)
Table IV,
Equilibrium Temperatures of the Standard and Light
_______
Standard Waters
No*
1,
2.
Bate
g
m
5/20/38
5/21/38
• #: ' 5/23/38
’ 5/25/38
011
Water
Equilibrium A T*
temperature
Saturated
with N
N
1.924
u
N
1.923
New oil sat.
with N
Same sample
N
Gr
1.921
1.931
+0.010
4 0.028
5,
5/23/38
n
ft
B
1.949
6,
5/23/38
!!
tt
N
1.928
7.
6/2/38
8.
6/2/38
New oil sat.
with N
Same sample
N
G
1.920
1.933
40.013
9,
6/Z-/38
U
tt
B
1.949
4*0.029
10,
6/3/38
if
n
N
1.926
11.
6/6/38
ti
tt
N
1.926
12.
6/6/38
U
tt
N
1.926
13.
*6/10/38
B
1.961
N
N
1.932
1.930
B
2.242
N
2.209
14,
15.
*6/10/38
6/10/38
16. **6/10/38
17. **6/10/38
New oil sat.
with N
New oil sat.
with N
Same oil
New oil sat.
with, B
New oil sat.
with N
+ 0.029
40.031
+ 0.033'
•»#
To insure the homogeneity of the saturated oil in
these two experiments 30 cc, were poured back and forth
between the two vessels several times and then divided
equally betweem them.
*-*.
In these two experiments the oil was redried,
twice distilled, treated as in
then saturated with
waters B and N respectively.
(19b)
*ts
o
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(20)
Solubility data are given in Table FT.
6.
Solubility of Water in the Oil
An approximation of the solubility of
water in the oil was obtained in the following manner!
A small drop of water was added to 200 cc. of the
dry oil and thoroughly shaken.
At the end of twelve
hours a single globule of water could be detected
in the oil on agitation, but at the end of eighteen
hours the droplets of water had completely- disappeared.
Another drop of water was then added and small drop­
lets could be seen when the oil was agitated at the
end of several days.
An exact quantitative determination
of the water was obtained as follows:
A 100 cc.
sample of the saturated oil was distilled under
reduced pressure at 50 - 60° C. through a drying
TJ-tube, fitted with
standard ground glass joints,
containing anhydrous magnesium perchlorate. ( Phos­
phorous pentoxide was tried as the drying agent but
was found to react with the oil.)
The oil was swept
out of the drying tube by air which passed through
concentrated sulfuric acid, two towers of soda-lime,
(20a)
Fable VT.
Solubility of the Oil in Water at Approximately
25° C.
ce#of water
in contact
with oil from
2- 4 days.
Grams of
AgCl (Av.)
Grams cf
oil per
100 c<r.
of water#
Fractional
part of oil
per part of
water.
100
0,004 9
0.004 1
0,00004!
Solubility of the Water in Oil at Approximately
25° C •
eo. of oil
saturated
with water#
Grams of
water per
100 cc#
of oil#
Grams of
water per
gram of
oi1 •
100
0.0277
0,000277
and a 30 X f inch pyrex tower packed with ??lass
wool and phosphors; , pentoxlde sealed directly to
the pyrex distilling flask.
The extra side arm
-
to the distilling flask through which the oil was
introduced was sealed during the distillation.
The ground glass Joints of the drying tube were
made perfectly air-tight by the use of stopcock
grease around the upperhalf of the Joint.
This was
a source of uncertainty and the second run was
carried out by sealing the top of the dry ground
glass Joints with picein cement.
At the end of
the experiment the picein cement was softened by
means of a small flame and completely removed by a
cloth dipped in benzene.
The method of weighing
the drying tube was similar to that described for
the determination of the density-temperature
coefficient ( see above p. 10) except that the
weighings were made on a balance accurate only to
0.2 mg.
The dry air was passed through the drying
tube until constant weights were obtained.
The two
runs checked each other within 0.2 m*?. Elsey
(21)
reports the
21. Elsey- Abstracts of papers, Phy.'and inorerAnic'
Division, Am. Chem. Soc.'meeting,Milwaukee, ^is
Sept. 5, 1938. Paper 64.
(22)
satisfactory determination of the water in freon
in a similar manner.
The data for the solubility
of water in the oil are given in Table VT.
Discussion of Data
In order to calculate the true density
of water containing an Unknown concentration of
deuterium oxide from
^T', the difference between
the equilibrium temperature of a drop of standard
water in the -oil saturated with standard water and
the equilibrium temperature of a drop of the unknown
water in the oil saturated with the unknown water,
and the net difference between the density-temperature
coefficients of the saturated oil (0.92 p.p.m. per
0*001° C*) and water (0.21 p.p.m. per 0.001° C.),
it is necessary that the change in density due to
the exchange of the water in the drop and that in
the saturated oil be less than the experimental errorr
and that the density-temperature coefficient of the
water saturated with oil " be,
equal to that of the
pure water.
In a 15 cc. sample of the oil there
(23)
is dissolved approximately 4.0 mg. of water, and
if we assume that the volume of
water*&i$eolve&> is
constant regardless of the deuterium content, it
would require a deuterium content of
37000,p#p*y*oto
change the density of the oil 1.0 p.p.m.
Therefore,
it seems reasonable to assume that the change in
density of the saturated oil with the deuterium
content of the saturating water does not cause a
measurable error as long as the deuterium concen­
tration is small.
If a 0.05 gram drop of B water, which
was 17*0 p.p.m* lighter than the standard watery
came into complete equilibrium with the 15 cc. of
oil containing 0.004' gm. of standard water, the
density of the drop
would become 15.74 p.p.m.
lighter than the standard water, and a 0.02 gm.
drop would become only 14*05 p.p.m. lighter than
the standard water.
^ile the size of the drops of
water were not accurately measured and varied slightly
for different experiments their approximate sizes
were from 0.02 to< ,0.05 gms.
The assumption that
there would be complete exchange between the drop of
water in the oil and the water dissolved in the oil
(24)
during a measurement may not be wholly Justified
since a large portion of the oil never comes in
contact with the drop of water, and the data illus­
trating typical measurements given in Tables II.
and VII*, and plotted in "Fig. 2 and ^*ig.S; show that
if an exchange did take place it was complete within
50 - 45 minutes*
The drop of water B, 17.0 p.p.m.
lighter than ordinary water, was introduced into
the oil at 10:00 o’clock and at 10:45 it was moving
through the oil at a uniform rate, (see Table VII.
and ^ig.S)
The straight line curve shows that the
distance of rise or fall of the drop through the oil
in unit time (3 minutes) continued to be a linear
function of the temperature for more than one hour,
which indicates that the density of the drop of
water was constant oVer that period and that the
change in density due to the exchange factor durinv
that period was less than the experimental error.
A complete exchange in this case would cause a change
of 2*0 p.p.m. in the density of a 0.05 gm. drop,
which is greater than the experimental error, since
a change of 0 .001 ° in temperature corresponds to
0.71 p.p.m. change in the density of the oil.
The
(24a)
Table VII.
Bata Illustrating a Typical Measurement
___________ Water B fY s-17.0 p.p.m.)
Time
Temperature
Cathetometer Readings
Initial
Final
Distance
cm. in
3 min.
10:45
1.946
29.06
29.23
-*-0.17
10:477
1.946
29.23
29.42
0.19
10:50
1.946
29.42
29.63
0.21
11:20
1.949
28.13
28.12
-0.01
11:23
1.949
28.12
28.12
0.00
11:26
1.949
28.12
28.12
0.00
11:29
1.949
28*12
28.12
0.00
11:39
1.951
28.08
27.97
-0.11
11:42
1.951
27.97
27.84
0.13
11:45
1*951
27.84
27.72
0.12
11:54
1.948
27.63
27.68
/•0.05
11:57
1.948
27.68
27.75
0.07
(24b)
40
30
>vo. yy\.
20
IO
20
30
40
1.946
1948
tSSO
1.95Z
1934
1.950
Fig. 5
Equilibrium temperature of B water in the saturated oil
(25)
data given in Table II* and plotted in Pier*2. far
G water show that the distance
of the rise or fall
of the drop of water was a linear function of the
temperature over a two hour period*
however, the change
In this instance,
in density Co. 8 p.p.m.) of a
0*03 gm. drop due to the exchange factor would be
of the same magnitude as the experimental error.
Furthermore, the ^ ^ values obtained in Table IV*
by subtracting the equilibrium temperature (2 *2 1 0 )#
for water N In 17 from the equilibrium temperature
(2*242) for water B in 16 where the influence of the
exchange factor was eliminated was larger by 0*004°
than the mean value of
iments.
A T in the previous exper­
This would correspond to a decrease in
density of 2.3 p.p.m., and tends to indicate that a
complete exchange between the drop of water and the
water in the oil had taken place.
Too much weight
should not be given to this experiment since it was
not checked by other experiments under similar
eonditions.
*
The equilibrium temperature increased with
redistillation of the oil indicating that the
oil was not absolutely pure.
(26)
The exchange factor offers an explan­
ation for the increase of the equilibrium tempera­
ture for the water II in the oil following the light
water measurements as seen in Table IV. 6 , 10, 11,
and 12*
However, there is some uncertainty here
since the exchange factor would not account for the
magnitude of the increase as noted, unless a drop
of the light water was lost On the sides of the
measuring vessel, -which was not the case in these
experiments#
IPurthermore, if the increase was due
to the exchange factor alone there should be a de­
crease of the equilibrium temperature of water N in
11 and 12#
Consequently, it seems necessary to
assume that part of the increase is due to the con­
tamination of the oil during the process of several
measurements.
Considering these measurements, it
is highly important that the procedure adopted in
measurements 16 and 17 of Table T7# be followed in
attempting to measure the density of water by this
method#
The densities of the standard light
waters (see Table V.,column 4 p19b) were calculated
by the equation:
(87)
v
AT rA
d
(
p
)
Y = ATb*^where
,i
O
k
dC
c
o
)*
7
^r-J
A T is 0.001° C. and
is the difference
between the equilibrium temperature of the standard
light water and that of the standard water in the
saturated oil*
If we substitute -0*92' p*p*m. for
Acf(o)
^ 7(see Table I* p. 12a) and -0*21 p.p.m. for
(from the International Critical Tables)
for an increase of 0.001^ C., the equation becomes
Y = AT[-o.7i/
The V
Values calculated in this manner are slightly
too high when compared with the true values as measured
by the quartz float method.
of
y
This numerical increase
is undoubtedly due to the invalid assumption
that the density-temperature coefficient of the
water saturated with oil is the same as that of
pure water.
If either of the unknown waters is
accepted as a standard and the decrease in density
considered as a linear function of the temperature,
the density of the other can be calculated from
AT
within 1 p.p.m. of its true density as shown in
Table V., columns 5 and 6 *
Therefore, the method
(28)
©an be used to calculate the deuterium content of
an unknown water, where the density is small, with
an accuracy of 0,001 $ deuterium oxide*
However,
it is not to be assumed that the linear relation
would hold over a wide density change since the
density difference in the waters compared in this
Investigation was only 10 p*p.m.
These experiments
and results are to be considered as preliminaries
and additional studies to be made in the near future
will be necessary to test the actual: merits of
the method*
With an accuracy of only 1 p*p*m* the
method as a routine laboratory procedure would have
no advantages over the improved
Falling Drop Method
as reported by Eeston, Rittenberg, and Schoenheimer (15)
*
However, the observations upon the solubility, density,
and volume factors, and consequently,the change in the
rate of rise or fall of the drop for the first three
or four minutes show definitely that the density
is not a direct inverse function of the rate of fall
of the drop, and further, that the deviation from
an inverse function is not due entirely to the dis­
tortion of the drop as a function of velocity but
to solubility factors as well.
Summary
A method of measuring the density of
a very small quantity of. water , which consists of
balancing a drop of the water in cyclohexyl chloride
saturated with the water being examined and comparing
its equilibrium temperature with the equilibrium
temperature of a drop of standard water in oil sat­
urated with standard water,has been studied*
By
standardizing the method with a water of known density
the density of a second water of small density dif­
ference can be calculated with an accuracy of 1 p.p.m.
The method is still in the experimental stage and
further studies will be necessary in order to prove
its real merits.
At present the method has no
advantages over the Falling Drop Method of Keston,
Rittenberg, and Schoenheimer since its accuracy is of
the same order as that reported by them.
(30)
I.
Isotopic Studies
A.
A Redetermination of the
Protium-Deuterium Ratio in
Normal Water
B.
The Microdetermination of the
Density of Small Quantities of Water
by the Balanced Drop Method .
1.
Urey, Harold C. and Gordon K. Teal
The Hydrogen Isotope of Atomic Weight Two .
Reviews of Modern Physics, 7, 34 (1935).
2.
Dufour, Comptes Rendus, 24, 1080 (1862).
3.
Richards, Theodore W. and John W. Shipley
A New Method for the Quantitative Analysis of
Solutions by Precise Thermometry.
Journal of the American Chemical Society
34, 599 (1912).
A Convenient Method for Calibrating Thermometers
by Means of Floating Equilibrium.
Journal of the American Chemical Society
36, 1 (1914).
4.
Richards, Theodopp W. and Gorham W. Harris
Further Study of Floating Equilibrium •
Journal of the American Chemical Society
38, 1000 (1916).
(31)
5.
Lamb, Arthur B. and R. Edwin Lee
The Densities of Certain Dilute Aqueous
Solutions by a New and Precise Method.
Journal of the American Chemical Society
35, 1666 (1913).
6.
Hall, Norris F. and Thomas 0. Jones
A Redetermination of the Protium-Deuterium
Ratio in Water.
Journal of the American Chemical Society
58, 1915 (1936).
7.
Gilfillan, E.S. and M. Polanyi
Mikropyknometer zur Bestimmung von
if
Verschiebungen im Isotopenverhaltnis des Wasser.
Zeitschrift fur Physika Lische Chemie
A 166, 254 (1935).
8.
Gilfillan, Edward Smith, Jr.
The Isotopic Composition of Sea Water,
Journal of the American Chemical Society
56, 406 (1934).
9.
Greene, Charles H. and Rodger J. Voskuyl
The Relative Proportions of Deuterium in Some
Natural Hydrogen Compounds.
Journal of the American Chemical Society
56, 1649 (1934).
(32)
10.
Rittenberg, D. and Rudolf Schoenheimer
Deuterium As An Indicator in the Study
of Intermediary Metabolism.
Journal of Biological Chemistry, 111, 169 (1935).
11.
Fromherz, Hans, Robert Sonderhoff and Heinz Thomas
Eine Einfache Methode zur Bestimmung des D2O Gehaltes kleiner Wassermengen •
Berichte der Deutschen Chemiechen Gesellschaft
70, 1219 (1937).
12.
Fenger-Erieksen, Karen, August Krogh and Hans Ussing
A Micro-Method for Accurate Determination of
D20 in Water.
Biochemical Journal, 30, 1264 (1936).
13.
Barbour, Henry G. and William F. Hamilton
The Falling Drop Method for Determining
Specific Gravity.
Journal of Biological Chemistry, 69, 625 (1926).
Blood Specific Gravity: It*s Significance and
A Hew Method for Itfs Determination.
American Journal of Physiology, 69, 654 (1924).
14*
Hochberg, S. and Viator K. La Mer
Microdetermination of Density by the Falling
Drop Method.
Industrial and Engineering Chemistry
Analytical Edition, 9, 291 (1937).
15.
Keston, Albert S., D. Bittenberg and Rudolf
Schoenheimer
Determination of Deuterium in Organic Compounds .
Journal of Biological Chemistry, 122, 227 (1937).
16.
Cri st, B. H., G .M . Murphy,afi&rHAroId Urejfrey
The Isotopic Analysis of Water.
Journal of the American Chemical Society
55, 5060 (1933).
The Use of the Interferometer in the Isotope
Analysis of Water.
Journal of Chemical Physics, 2, 112 (1934).
17.
Norris, J.F.
n-Butyl Chloride
CHq (CHg )gCHgOE + HC1 + (ZnCl2 )■ CH3 (CH2 )gCHgCl + H2(D
Organic Syntheses, John Wiley and Sons, New York.
Collective 1, 137 (1932).
(34)
18.
Gucker, Frank T*. Jr., Fred W. Gage and Charles
E Moser
The Densities of Aqueous Solutions of Urea At
o
o
25 and 30 and the Apparent Molal Volume of Urea.
Journal of the American Chemical Society
60, 2582 (1938).
19.
Richards, Theodore William
A Method of Standardizing Weights •
Journal of the American Chemical Society
22, 144 (1900).
20.
Thiessen, G.W. and L.J. Frost
Constant-Temperature Bath Employing Thermionic
Control *
Journal of Chemical Education, 12, 72 (1935).
(1)
A Study of the Properties of
Lithium Glass Electrodes
Review of Literature
In a very careful study of the effect
of the chemical composition of glass upon its
behavior as a hydrogen electrode, Maclnnes and D o l e ^
found that the composition corresponding to the low(2 ) triangular
est temperature found on Morey’s
melting point diagram for the system, CaO -SiOg -Na^O
( CaO 6$, SiOg 72% 9 TTagO 22%), was the best for
determining the hydrogen-ion activity.
Electrodes
made from thin membranes of this glass gave the
lowest asymmetric potential, the lowest electrical
resistance, and the smallest error in alkalin©
solutions of any glass studied.
These observations
are in agreement with those of other investigators
upon the effect of the chemical composition of the
1*
2.
Maclnnes aid Tde, J. Am. Chem. Soc., 52,29(1930)
Morey, J. Soc. Glass Tech., 9, 232 (1925)
I.C.T. M Graw-Hill Book Co., Vol. 11,p. 97
(S)
glass upon its electrode properties.
the assumption that lithium,
smaller
(3 4 &5)
9 9
Upon
with an atomic volume
than that of sodium, when substituted wholly
or in part for the latter in the composition of
glass, would
yield a product, which as electrodes,
would eliminate some of the error in alkaline
solutions, Maclnnes and Bole studied glasses of the
following compositions:
1.
SiOg
72 wt . Percent
CaO
6
Li2°
2.
SiOg
72 wt . Percent
tt
ti
CaO
6
22 "
tt
LigO
NSgO
SiOg
72 "
if
CaO
6
■
Li20
4
NagO
tt
tt
2
tr
tt
20
it
it
SiOg
72
w
ft
tt
CaO
6
»
tf
"
it
LigO
11
it
If
18 ”
tt
KgO
11
rt
tf
4.
Electrodes made from membranes of glass No .
1 .
had
a low resistance, low asymmetric potential, but the
potentials were in error even in solutions of pH 8
on the second day and the error, which was about
Hushes, J. Chem. Soc., 491, (1928)
I; Eller and Wright, Proc. Nat. Acd.of Sei. 14,936(1928)
5 . Kahler and DeEds, J. Am. Chem. Soc.,53, 2998 (1931)
3
(3)
the same as that of electrodes made from the best
sodium glass, increased rapidly in strong alkaline
solutions after the first day*
All electrodes made
from membranes of glasses 2 , 3, and 4 had
very high
resistances and large errors, both of which increased
rapidly with time in akaline solutions*
In their
ability to function as hydrogen electrodes, none of
the lithium glasses were found superior to the best
sodium glass, which is now commonly referred to in
the literature as Corning 015 glass.
and Passynsky
However, Ssokolov
(6 ) have recently reported that elect­
rodes made from glass having the composition, SiOg,
80$; CaO, 10$; and LigO, 10$, function as r perfect
hydrogen electrodes in solutions with a pH as
high as 12*5#
Statement of Problem
Although Maclnnes and Bole did not
investigate a glass of the identical composition as
that used by Ssokolov and Passynsky, it was felt
6d#
Ssokolov and Passynsky, Z. physik; Chem. A 150,
366 (1932)
(4)
that the conclusions of the two independent inves­
tigations were sufficiently contradictory to warrant
a reinvestigation of the electrode properties of
glass of the exact composition recommended by them.
Furthermore, the tremendous increase in the use of
the glass electrode in the last few years makes
the finding of a glass that would practically elimi­
nate the error of this electrode in alkaline
solutions of paramount importance.
Some Preliminary Experiments
Error of Sodium Glass Electrodes at 50° C.
1.
Apparatus and Materials
In a previous investigation Amis and
Gabbard
(7)
used the vacuum tube circuit described
by DuBridge and Brown
(S)
to build a vacuum tube
potentiometer very satisfactory for sodium glass
electrode work, and it was decided to use the same
potentiometer in this investigation.
shown in Fig. I*
7.
8*
The circuit is
The entire potentiometer,except
Amis and Gabbard, J. Am. Chem. Soc., 59, 557(1937)
DuBridge and Brown, Rev. Sci. Instru., 4,532(1933)
^
v
r-V\AA~
-AAA/'
c?
(4a)
c?
a
AAAA
J V W N A -
Fip. 1. Potentiometer Circuit
(See Key p. 4b)
v w s vyv
(4b)
Key to Fig* 1
B~ five 3 volt, 300 ampere hour dry cells-*
F- a l/8 ampere fuse*
G- an L & IT type R Ko* 2500-e galvanometer*
I- a 0 - 100 range millimeter*
K^— a single contact , short circuiting key*
Kg-* a single contact tapping key for protecting
the galvanometer*
R*^ and. Rg are 50 ohm potentiometers.
R3 , R4, and R5 are respectively 50, 50, and 10,000
ohm rheostats*
Rg, R 7 , R8, and Rg are respectively 6000, 2000,
10,000 and 300 ohm resistors.
Sg is a multiple contact double pole double throw
switch for changing the galvanometer from the
plate circuit to the type K potentiometer
circuit and vice versa.
S4 is a single pole two position switch.
S3 , Sg, and Sg are multiple contact single pole
single throw switches*
V is a l-T range voltmeter*
K-Pot is an L & N type K potentiometer*
$R-£>4 is a General Electric FP-54 vacuum tube.
the galvanometer and the tapping keys, was enclosed
in a galvanized iron case, from the outside of which
the controls of the Leeds and Northrup Type K poten­
tiometer were operated by means of brass extensions.
The leads to the galvanometer and the tapping keys
were of microphone wire*
The parts of the Dutridge
and Brown circuit were mounted on high grade bakelite fastened to the inside of the iron case with
7 cm. beehive porcelain insulators, and were operated
by bakelite rod extensions. Special attention was
given to the shielding and insulation of the FF-54
pllatron tube
and the grid switch.
These were sup­
ported by bakelite plates mounted on beehive
por­
celain insulators and enclosed in a small galvanized
iron case, within the larger iron case, made air­
tight and kept dry by means of phosphorus pentoxide.
The switch,which was connected to the control errid
of the vacuum tube* consisted of a copper wire bent
at right angles at each end and dipped into mercury
cups joined respectively to the ground and the cell ,
was
controlled by a bakelite rod extension.
The
leads to the switch and the tube entered the small
case through bakelite by means of porcelain insulated
(6)
connectors*
For the sake of developing a certain
amount of technique in measuring the error of the
glass electrode in alkalime solutions and Turtheh to
satisfy ourselves as to the dependability of the
potentiometer for this type of work, the error of
the sodium glass electrode in alkaline solutions at
30° C* was measured*
The general experimental procedure
(9 )
first introduced by Dole
of comparing the glass
electrode directly with a hydrogen electrode in the
same solution over a wide pH range while holding the
concentration of the metallic ion constant was
followed*
The comparison cell shown in Pig* 2
was suspended in an oil bath whose temperature was
controlled within ± 0*01° C. by use of a mercury
regulator operated by means of an electrical relay*
The oil in the bath was thoroughly agitated with an
air driven motor, and the entire bath was inclosed
in the large metal box while the measurements were
9.
Dole,
Am* Chem. Soc*, 53, 260 (1931)
(Ga)
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being made. The glass electrodes were of the Maclnnes(10 )
Bole
type made by sticking the thin membrane on
to the end of a hot soft glass stem by means of a
water aspirator*
A water trap in the suction line
controlled the pressure exerted upon the glass
membranes*
A silver-silver chloride electrode of
(11}
the reduced oxide type
was placed in 0*1 N.
hydrochloric acid inside the glass electrode and served
as a reference electrode*
The hydrogen electrodes
were made in the usual manner, and were cleaned and
replatinised after each series of experiments* The
i
calomel electrode was made from Leeds and Northrup
chemicals, specially purified for this purpose*
Two glass and two hydrogen electrodes were used in
the experiments*
The two hydrogen electrodes were
checked against each other from time to time and
were shorted with each other during the actual read­
ings*
Commercial tank hydrogen was purified by
passing it through
'concentrated .potassium hydrox­
ide, through distilled water, over a copper catalyst
10*
Maclnnes and Bole, Ind* Eng* Chem* Anal. Ed.,
1, 57, (1929)
1 1 * Jones and Hartman, J. Am. Chem. Soc.,37, 752(1915)
(8 )
heated to 450 - 500° 0. and washed by the solution
upon which the measurements were being made before
entering the comparison cell*
Part of the hydrogen
was diverted from the main chain ©f apparatus to
saturate the sodium hydroxide which was used to in­
crease the pH of the solution*
Rubber connections
between the catalyst and the cell were avoided with
one exception.
In this case the connection was made
with thick pure rubber impregnated with paraffin.
The concentrated stock S°!utions of
sodium acetate and sodium hydroxide were made from
the dry C.P. chemicals and the sodium determined
analytically as sodium sulfate.
The weaker solutions
were made by diluting the stock solutions in a
calibrated flask.
2.
■Experimental Proce dure
With glass, hydrogen, and calomel
electrodes immersed in the same solution the following
cells could be formulated.
H yd r o gen-Glass
Pt, Hg/soln/ glass/ 0.1 H. HCl/
AgCl/fig
B.
Hydro gen-Caloifl&l
Pt, Hg/soln// sat .’XClAlpClATfT
(9)
C.
Glass-Calomel
A g A g G l / )0.1 H. HCl/ glass/ soln.//
sat. KOI/ HgCl/ Hg
•
f
t
*
The values of cells B. and a* were
obtained at various pH values ranging from 7 to 12
for 0.1 and 1*0 H* sodium acetate solutions*
The
value of cell A. was obtained by subtracting the
value of cell C. from the value of cell B.
The pH
of the solution wap increased by adding sodium
hydroxide of the same sodium-ion concentration -• •
saturated with the purified hydrogen.
It can be
seen from observing cell A. that as long as the glass
electrode is acting as a perfect hydrogen electrode
the
value of the cell should be constant.
The constant
for the cell was obtained by measuring the value
of the cell by difference at a pH of 8.5, or below
where it is known there is no glass electrode error.
The
value of the glass electrode error, (A£)>
is
the
deviation of the value of the cell A. from the
cell constant.
The pi of the solution was calculated
from the value of cell B. by means of the following
equation:
pH _
~
Eg - Eq -f Ep
0.0601
(10)
Eg is the value of cell B , Eg is 0.2420vf the value
of the saturated calomel cell at 30° C., and E- is
the correction necessary to bring the hydrogen
pressure to 760 mm. of mercury.
In Table I. is an
illustration of the data as taken in a typical
experiment •
liie errors of the sodium glass
electrode at 30° 0. in sodium-ion concentrations of
0.1 and 1.0 H* are given in Table II., and Hig. 3
shows a comparison of the data with the data of Dole
and Wiener in solutions
of the same concentrations
**
....
m
at 25°
and 50° C..
It is evident from Table I.
that the results of this investigation are extremely
concordant since the difference between two glass
electrodes never was as much as one
millivolt.•
•
The average difference was only 0.3 millivolt.
The data was accepted as conclusive proof pf the
accuracy of the described potentiometer for glass
electrode work.
Materials and Experimental
Procedure for the Study of the Lithium
Glass Electrode
(10a)
80
TO
6O
50
30
20
PH
Fiff. 5. Error of the plass electrode at
compared
with the data of Dole and wiener^14' at 25 and
50° C.
14. Dole and Wiener, Electrochemical Society, 72, 25 (1957)
(10 b)
(IOC)
Table XI
Error of the Glassr Electrode at 30° C.
Solution______
0.1 N* Na________
Electrode
pH
Error* in mv.
No.l
o.l N. Na
Electrode
iM
No.2
Error in imr.
No.l
NO,
9*26
3.0
3.4
10.17
13*0
13.4
9.65
4.5
4.9
10.85
25.8
26.5
10*05
6.7
6.9
12.22
71.7
71.9
10*66
10.8
10.1
12*54
87*6
87.5
11*27
19.1
19.5
12.71
97.1
97.0
11*99
35.2
35.5
1*
Preparation of the Electrode®
The Corning Glass Works very kindly
prepared for us in 1934 a glass, in the form of
opalescent and opaque rod® to which they assigned the
No* 121 B., composed of
10$.
SiOg 80$; CaO 10$; and kigO
This i® the identical composition of the glass
reported by the Russian workers.
The Coming Class
Works informed us that the glass would not be very
workable and would be subject to devitrification,
which was verified in practice.
The preparation of
the Maclnnes- bole type of electrode from this glass
was found to be extremely difficult*
When the glass
was gently heated it became yellowish, and additional
heat produced a relatively soft, clear glass of a
very low viscosity*
All attempts to blow the glass
into membranes at this temperature failed, and the
glass would rupture producing what might be referred
to as a ,!blowout.u bevitrification and hardening was
also present at this stage*
buring the application
of sufficient heat to keep the glass soft the color
changed to dull red.
At a higher temperature, which
required the use of an oxygen flame, bubbles appeared
which in turn disappeared at a still higher temperature
giving again
a relatively easy flowing, clear glass.
It seems that Ssokolov and Passynsky
f6v
* made their
bulb type electrodes from this final clear glass, but
in this research, it was found that the best glass
films could be blown from the glass in the dull red
stage.
A 2 mm. tube of ordinary soft sodium glass
With a piece of rubber tubing attached for blowing
was clamped in a vertical position.
of the lithium glass was placed
A small piece
upon the end of the
tube and heated with extreme care by the use of an
oxygen hahd torch#
When the lithium glass reached
the dull red stage, it was blown out into thin
Iridescent films from which the electrodes were made
as described for those of sodium glass.
A higher
temperature was necessary to seal the films on to
the ordinary sodium glass stems, than was required
f
in making the ordinary glass electrodes, and only a
small percent of the electrodes were of any value.
This was attribu,ted to
the dIffiou.11y of ge11 in?
the lithium glass membrane to make a perfect seal
with the ordinary sodium glass.
Electrodes were
also made from film blown from the glass in the
final clear stage.
A small amount of approximately
(13)
0*1 N. hydrochloric acid was placed inside the elec­
trodes which were allowed to stand in approximately
0 *1 .N* acid for at least twenty-four hours before use.
. 2>* Resistance of Lithium Glass Electrodes
The resistances of the electrodes were
measured by the vacuum tube potentiometer by employ­
ing the input circuit illustrated in Fig*4 *
^ith
one volt constant potential on the potentiometer (K.P.)
switch S. in position 1 and switch. S3 closed, there is
a certain drop of potential across the glass electrode,
(Rg), giving a fixed grid potential and resulting
in a definite plate current, indicated by the galvan­
ometer as the null point.
and Sl at position 2>
Now with
open, Sg closed,
Rv can be varied until the
galvanometer is brought to the null point, and the
resistance of the electrodes can be calculated from
the following equations:
Rl t
(1)
(2 )
f3 )
(Rv -f-Rg) Iq - CRi+ Rg) 5, or Ig/rf
RgXg = Rgli> or
Rv t
Rg/R2 =- IgAl
From (1 ) by substitution
^1 ^2 ^ ^2 Rg
RgRv 4* Rg R g —
Rp
ra
ro
r—!
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electrodes. H- and T?2" raQ£Ohm resistors, R«- variable
resistor, 1- 10,000 obms. (See Fie1. 1, p. 4a for
remainder of circuit)
(13a)
There fore, Rg =g R^ Rg/ Rv
The cell (see T'ig* 4) used for measuring
the resistances:: of the electrodes was the same as
that used for the asymmetric potential and consisted
of a battery Jar
5 in. X 6 in. X 3.5 in. with the
silver-silver chloride and the glass electrodes
supported by a plate of high grade bakelite.
electrodes were S.5 in.
The
apart and at least 1*5
in. from the sides of the glass vessel which rested
on a glass plate.
The electrodes dipped into approx­
imately 0.1 R. hydrochloric acid, the same concen­
tration as the acid inside the glass electrode.
The
accuracy of the resistance measurements was easily
within 500 megohms.
The results' fbfc the resistance
measurements are given in Table III.
After the electrodes had been packed
in cotton fifteen months, one of the electrodes, the
■onlyone functioning, was found to have a resistance
of the magnitude of approximately 4000 megohms.
There was some uncertainty as to the exact magnitude
of the resistance since it seemed to increase slowly
during the measuring process'.
This increase was
attributed to polarisation of the electrode.
The
(14a)
m
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of Lithium
Glass Electrodes
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(IS)
resistance was measured as previously described, and
also by measuring the drop of potential with the
vacuum tube potentiometer betrween the terminals of
the silver-silver chloride ele ctrodes with the passage
of a known direct current*
The minimum resistance
was 3500 megohms while the maximum res istance was
9000 megohms*
It should be stated that this large
change in resistance during the actual time of
measurement was not noted in the earlier measure­
ments nor was it noted in the case of a sodium glass
electrode of approximately 1000 megohms resistance*
3*
The Vacuum Tube Potentiometer for the
Highly
Resistant Lithium Glass Electrode Measurements
When an attempt was made to measure
the asymmetric potential of these electrodes by
connecting tirent directly to the grid of the potentio­
meter* as was done in the case of the sodium glass
electrodes, it was found that their resistances
were
of such magnitude that the plate current slowly
decreased* indicating that electrons were slowly col­
lecting on the grid of the tube*
The decrease was
alow but of a uniform rate, and, if given enough time,
(16)
would finally, after several minutes, reach a
minimum.
If the switch or leads' to the cell were
grounded at this point, the charge would’ apparently
leak off and gradually build up again to a maximum
value.
The BuBridge and Brown circuit in the
potentiometer
fsee Fig. 1.) was
modified to make
use of the Impulse Amplification principle employed
on many commercial pH meters, and fully described by
(12 )
Ellis and Kiehl#
A 150 micro-microfarad isolantite air condenser was introduced into the DuBridge
and Brown circuit as shown in Fig. 5. The condenser
-13
would require 1.5 X 10
coulombs per millivolt
charge.
When S is at position,©, the condenser is
charged to ground potential and thus the grid is at
ground potential which gives a definite plate current,
and any suitable reading on the galvanometer can be
selected as the null point.
If S is now switched
to C, the condenser is charged to the unbalanced e.m.f.
of the cell producing a change of grid potential and
thus a change in the plate current.
When the e.m.f.
from the potentiometer K exactly balances the e.m.f.
12.
Ellis and Kiehl, Rev. Sci. Instru.,4, 131 (1933)
/;+e.
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electrodes.
C-a 150 mtc'romiorofarad
p. 4a for remainder of circuit)
•p
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Ph
of very high resistance
condenser.
(See Fig.l,
(16a)
10
fcl
•H
Ptn
(17)
of the cell, position C is also at ground
potential
and there will be no change^ in the plate current
when S is moved from C to 6 , and the reading on the
potentiometer is the e.m.f* of the cell*
When the
potentiometer as described was tried out, it was
found that the plate current decreased as before
except more rapidly even when the condenser was at
ground potential*
'This was due to the charge on the
grid caused by small grid current and -fee very high
16
grid resistance (10
ohms) of the tube* This
difficulty
was avoided by providing- a leak to the
ground for the charge by means of a 104,400 megohm
resistor.
The ability of the potentiometer to measure
the potential of high resistant cells was then tested
by measuring the potential of a standard cell in
series with high resistors*
The results were very
satisfactory and are given in Table IT*
4*
The Asymmetric Potentials of the Lithium Glass
Electrodes
The measuring cell as described (see
above p*14 ) was -Used to measure the asymmetric
potentials*
With approximately 0,1 H, hydrochloric
acid in the measuring vessel and In the lithium
(17a)
Table IV.
Vacutun-TiEbe Potentiometer Measurements on a
Standard Cell with Known Resistances inserted
in Series
Ho Extra
1220 Megohms
2940 Megohms
Resistance
in Series
in Series
Volts
Volts
Volts
1.01500
1.01507
1.01483
1.01503
1.01480'
1.01497
1.01505
1.01485
1.01500
1.01487
1.01502
glass electrode the potential between the two silver
silver chloride electrodes should be zero, and any
deviation from1 this vralue is due to the asymmetric
potential of the glass films of the glass electrodes
Despite the fact that the ele ctrodes had a^ed in
dilute hydrochloric acid for three months, the asym­
metric potentials were large and w ould not come to
a constant value over a long period of time.
Some
typical data are shown In Table V.
Electrode Ho. 4 was made by heating
the -molten glass to a higher temperature according
to the procedure of Ssokolov and Passynsky.
The
electrodes made at this higher temperature showed a
higher asymmetric potential and also higher resist­
ances than the electrodes made from the glass at
lower temperatures.
5.
The Sensitivity of the Lithium Glass Electrode
to pH Change
Electrodes No. 1,
3, and 5 were
placed in the measuring cell in a buffer solution
of pH
8 #79,
and the potential of the cell C was
measured over a period of time.
The electrodes
(18a)
Table V.
Asymmetry Potentials of Lithium Glass Electrodes
No.l June 9
•Tim©
volt
5:05 p.m. 0.0555
Bb.if June 10
Time
Volt
1:45 p.m . 0.0363
No. 4 June 10
Time
Volt
9:45 a.m.0.0563
3:30
0.0406 22: 45
0.0320
11:00
0.0609
3:55
0.0303
3:30
0.0304
11:40
0;055V
4s20
0.0265
4:35
O'.0280
4*50
0.0268
5:40
0.0253
9:00
0.0216
9:30
0.0215
10:30
0.0225
1 0 ;00 a.m. 0.0319
(19)
were transferred to a buffer solution of pH 4.06 and
the potentials of the cell again measured.
A sodium
glass electrode showed the correct change in poten­
tial between the two buffer solutions to be 0.2813v.
Th© data for the lithium glass electrodes are shown
in Table VI, and it can be seen that these electrodes
fail V>ery badly to function as hydrogen electrodes
even when the maximum change in potential is consid­
ered.
In addition to giving an ever changing poten­
tial no two electrodes gave the same potential, and
only one electrode gave a change of potential
within four hundredths of a volt of the correct
thermodynamic change in potential.
It was felt that the change
In potential
might be due to an exchange reaction between the sodium
ions in the solution and the lithium ions in the glass.
Three of the electrodes including No. 3 were placed
in a 0.2 N. lithium acetate solution and the potentials
of cell C
measured at intervals.
However, there was
no improvement in the behavior of the electrodes
in this solution.
The stem of the
No. 3 electrode
and the entire measuring cell were paraffined care­
fully and the potentials ai?ain measured at intervals
(19a)
CD
fr-
o
>
C
25r
at Approximately
Cell e.m.f.'a
Saturated Calomel-Glass Electrode
Table V I .
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(20)
over a two hour period, but the results were about
the same as in the previous experiments*
Since the lithium glass electrodes had
large varying asymmetric* potentials and since they
failed completely even in the neutral range to function
as hydrogen eleetrodes, no attempt was made to try
their behavior in alkaline solutions.
The measure­
ments could not possibly have any significance.
6*
A Comparison of the Behavior of Lithium Class
Electrodes with the Behavior of Sodium Class
Electrodes of Approximately the Same Resistance
in the Same Solutions
The question had been raised as to
the possibility of the abnormal behavior of-the
lithium glass electrodes as shown in the preceding
experiments being due to the inadequacy of the insu­
lation and shielding of the potentiometer in spite
of the fact that it proved very adequate in measuring
the potential of the standard cell in series with
the high resistors as shown in Table IV.
It was argued that in measuring cells
where there is a much wider distributed capacitance
(SI)
there is? a much larger
static pick-up.
susceptibility to electro­
In order to answer this question
sodium glass electrodes of films thick enough to
have a resistance of the same order of magnitude as
the lithium glass electrodes were made and their
behavior compared with that of the lithium glass
electrodes on the same potentiometer.
It was assumed
that if the high resistant sodium glass electrodes
gave constant asymmetric potentials and constant
potentials when measured in buffered solutions,
then the abnormal behavior of the lithium glass
electrodes could be attributed to the lithium glass.
The data in Table VII. show the results of these
experiments*
It can be seen that while the sodium
glass ele ctrodes were somewhat sluggish and slow
in reaching a constant potential in both the
asymmetric potential measurements and the pE measure­
ments, they did come to a constant potential in
i
every case, which is not- true of the lithium glass
electrodes in a single case,
Therefore,
it is
felt that the abnormal behavior of the lithium
glass electrodes can be attributed directly to the
lithium glass itself.
The unpublished results of
(21a)
Table VII,
Data Comparing the sehavior of Lithium Glass
Electrodes with Sodium ,flass Electrodes of High
Resistance
Potential Measurements of Standard Cell in Series
with High Resistance
Time
12:00
2:3-
Ho extra resistance
1,0155 volts
1,0155 *
Tjme
2:40
3:00
1550 megohms
1,0155 volts
1,0155 n
Asymmetric Potentials
Ha— 16
1250 megohms
Time
volts
.0143
3:05
,0142
3:20
,0155
5:45
.0158
4:10
,0126
4:45
,0123
5:20
.0102
8:30
9:05t>.m..0102
9 :45a!#tn«0102
10:20
.0102
Ha— 18
3300 megohms
Time
volts
3 :00
3:25
3:50
4:17
4:56
5:45
.0123
.0096
.0086
.0077
.0073
.0073
Li— 1
4000 meerohms
Time
volts
5:55
6:25
7:45
.1050
.1071
.1123
.1082
8:10
.1056
8:45
9:03 p.m. .1045
7:55 a.m. .0965
.0947
8:30
.0965
9:05
.0890
9:55
(21b)
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Table
VII.
Continued
Calomel-Glass
Electrode
Measurements
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(22)
The National Technical Laboratories, Pasadena,
California,
with glass of the composition given by
• *
Ssokolov and Passynsky also .justify this conclusion.
The true change in potential, ( & £ ) , between the
two buffer solutions of different pH ^alue was shown
by the value of
£ of the sodium glass electrode of
low resistance to be 0.3394v*
The
deviation
error in case of the high resistant sodium glass
electrodes amounted to 0 .02v
in one instance and
0*009v in the other two electrodes, while the minimum
error for the lithium glass electrodes amounted to
more than one tenth of a volt.
The deviation
of the sodium electrodes from the true
r,
change
raises the question as to the relation of the thick­
ness and resistance of the glass electrode to its
ability to follow the true pH change.
‘it has been observed by many inves(3,13,5)
tig©tors
that thin electrodes and consequently
low resistant electrodes ffollow the pH change of the
solution more closely than high resistant electrodes.
*
13.
Private communication'
M0 rton, J. Chem. Soc., 256, (1934)
(23)
In the ease of high resistant electrodes made from
tubes of soft glass Kahler and DeEds1 found that the
deviation error
was due to the difference in the
depth of immersion of the
trodes*
two surfaces of the elec­
In such cases a part of the current which
flows through the electrode must first flow up a
conducting film of solution on the surface of the
glass*
This results in a potential between the
film and the liquid different from that between the
two liquid surfaces since the film would have
different concentration from the solution.
a
The
deviation was designated by them as the11deviation
film11
and can be eliminated in this type of electrode
by making the levels of the liqiud inside and outside
the electrode the same.
Since the depth of immersion
was approximately the same in these experiments, the
ttdeviation film"
error.
does not explain the
AS
deviation
'furthermore, in the Maclnnes-Dole type of
electrode where the resistance of the supporting stem
is many times that of the film, the”depth of immersion”
error, ortfdeviation film” , has not been observed.
The idea of electrical leakage, thoroughly discussed
by M o r t o n s e e m s the more reasonable explanation
in these experiments.
Since the resistances of
O'
the electrodes were of -the order* of 1 X 10
ohms,
the parallel leakage resistance would necessarily
have to be of the order of 10
12
ohms, and it is
not difficult to see that leakage can occur readily
under such conditions.
It is difficult
to understand
however, why the sodium electrode of the lower resist­
ance should have the
larger'A £ deviation error.
It has been observed by the authors in many instances
that sodium glass electrodes of relatively low
resistances
failed to follow the pH change in the
neutral or acid region where the electrodes are
ordinarily expected to function as perfect hydrogen
electrodes,
and for this reason it is felt that a
thorough study of this problem is desirable, and
will be undertaken- In the near future.
Conclusions
It is concluded that the lithium glass
electrodes of the composition recommended by Ssokolov
and. Passynsky are not superior to those made from
Corning 015 glass.
In fact,
it is believed that
they do not function as true hydrogen electrodes in
in any pH range and have no advantages over the sodium
glass electrodes*
No explanation has been found
for the disagreement between the results of this
investigation and those published by the Russian
workers unless they failed to report the correct
composition of their glass*
II.
A Study ojf the Properties of
Lithium Glass Electrodes
X.
Maclnnes, D.A. and Malcolm Dole
The Behavior of Glass Electrodas of Different
Compositions,
Journal of the American Chemical Society
52, 29 (1930).
2.
Morey, Journal Society of Glass Technology
9, 232 (1925).
x
International Critical Tables
McGraw-Hill Book Company, New York, New York.
2 , 97 (1927).
3.
Hughes, Walter Scott
On Haber's Glass Cell#
Journal of the Chemical Society, 491 (1928).
4. Elder, L.W. Jr. and W.H. Wright
pH Measurement with the Glass Electrode and
vacuum Tube Potentiometer.
Proceedings of the National Academy of Sciences
14, 936 (1928).
5.
Kahler, H. and Floyd DeEds
The Glass Electrode.
The Study of Various Characteristc
Journal of the American Chemical Society
53, 2998 (1931).
(27)
6.
Ssokolof, s.I. and A.H. Passynsky
II
Uber Glaselektroden*
Zeitschrift fur Physika Lische Chemie
A 160, 366 (1932).
?•
Amis, E.S. and J.L. Gabbard
A Comparison of Hydrogen, Quinhydrone and Glass
Electrodes in Magnesium Sulfate Solutions.
Journal of the American Chemical Society
59, 557 (1937).
8.
Du Bridge, Lee A. and Hart Brown
An Improved d. c. Amplifying Circuit.
Review of Scientific Instruments, 4, 522 (1933).
9.
Dole, Malcolm
Glass Electrode Measurements by Means of a Galvanometer
with Condenser Attachment.
Journal of the American Chemical Society
53, 260 (1931).
10.
Maclnnes, Duncan A. and Malcolm Dole
Tests of a Hew Type of Glass Electrode.
Industrial and Engineering Chemistry
Analytical Edition, 1, 57 (1929).
(23)
XI.
Jones, Grinnell and Miner Louis Hartmann
Properties of Silver Iodide Interpreted in Relation
to Recent Thermodynamic Conceptions.
Journal of the American Chemical Society
37, 752 (1915).
12.
Ellis, Samuel B. and Samuel J. Kiehl
A Practical Vacuum-Tube Circuit for the Measurement
of Electromotive Forces.
Review of Scientific Instruments
4, 131 (1933).
13.
Morton, Charles
The Effect of Electrical Leakage on the Electro­
motive Behavior of the Glass Electrode.
Journal of the Chemical Society, 256 (1934).
14. Dole, Malcolm, and B. Z. Wiener
The Theory of the Glass Electrode. IV. ^emoerature
Studies of the Glass Electrode Error.
The Electrodhemical Society 72, 25 (1937)
Vita
The write* of this thesis was horn
June 2, 1902 in Owsley County, Kentucky, and
received his elementary education in the public
schools of that county#
He attended the Berea
Hormal School", Berea College, Berea, Kentucky,
and received
a
diploma from that school in 1©B1#
His undergraduate work was done
at
Berea College
and at the University of Kentucky, Lexington,
Kentucky#
He received the Bachelor of Science
degree from the latter institution in 1926#
He taught two years in the elementary schools of
Harlan County, Kentucky* taught elementary biology
and chemistry for one year in the city schools of
Harrodsburg, Kentucky? was principal of the Dixie
High School in Henderson County, Kentucky for two
years? and went to the university of Kentucky in
1928 as an instructor in chemistry#
He received
the Master of Science degree from the university
of Kentucky in 1932 and the Doctor of Philosophy
degree from northwestern University in 1940#
James Lawrence Gabbard
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