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Scholars' Mine Scholars' Mine
Masters Theses Student Theses and Dissertations
1950
The determination of the solubility of hafnium oxide in aqueous The determination of the solubility of hafnium oxide in aqueous
solution by the radioactive tracer technique solution by the radioactive tracer technique
Hampden O. Banks
Follow this and additional works at: https://scholarsmine.mst.edu/masters_theses
Part of the Chemistry Commons
Department: Department:
Recommended Citation Recommended Citation Banks, Hampden O., "The determination of the solubility of hafnium oxide in aqueous solution by the radioactive tracer technique" (1950). Masters Theses. 4822. https://scholarsmine.mst.edu/masters_theses/4822
This thesis is brought to you by Scholars' Mine, a service of the Missouri S&T Library and Learning Resources. This work is protected by U. S. Copyright Law. Unauthorized use including reproduction for redistribution requires the permission of the copyright holder. For more information, please contact [email protected].
THE DETERMINATION OF THE SOLUBILITY OF HAFNIUU OXIDE
IN AQUEOUS SOLUTION BY THE RADIOACTIVE TRACER TECHNIQUE
BY
H~APDEN O. BM~KS, JR.
A
THESIS
submitted to the faculty of the
SCHOOL OF r,IINES AND HETALLURGY OF THE UNIVERSITY OF HISSOURI
in partial fulfillment of the work required for the
degree of
MASTER OF SCIENCE IN C~1ISTRY
Rolla, Missouri
1950
ii
AC KNOWLEDGEIJENT
The author wishes to express his sincere appreciation
to Dr. R.A. Cooley, Associate Professor of the Department
of Chemical Engineering, lJissouri School of Mines and
Metallurgy, for his invaluable suggestions and guidance in
the preparation of this thesis.
TABLE OF CONTENTS
PageAoknowledgement •••••••••••••..••••••••• 11
ill
L1st of 1llustrat1ons••• ......... . ••. .• 1v
List of tables •••••.•..••••••••••••••••• v
Introduction••• ••••••••••••••••••• 4 •• • •• 1
.................Review of literature. ..5
Equ1pment 7
Chemicals used•••••••••••••••••••••••••• 9
Exper1mental work••••••••••••••••••••••12
Validity of results •••••••••••••••••••• 37
Energy of Solution••••••••••••••••••••• 38
Debye-Huckel theory•••••••••.•••••••••• 39
Solubility Product ••••.• • • • • • • • • • .••••• 41
5l.1lnlnaI"y•••••••••••••••••••••• , ~ ~ , •••••• 44
SugGestions for further study•••••••••• 45
Appendix••......•••.....••••...•••••.•• 47
Bibl1ography••••.•.••••••...••••••••••• 48
Vitat ....••..........................•• 49
iv
LIST OF ILLUSTRATIONS
Fig. Page
1. Neutron bombe~dment of hafnium 180 atom ••••••••••••10
2, Lauritsen Quartz Fiber E1ectroscope••••.....•••••••13
3, Sensitivity 9lot for Lauritsen Electroscope•••••••• 14
4. Sensitivity plot for counter tube•••••••••.••.••••• 16
5, Horizonte~ counter tUbe supoort •••••••••••••••••••• 17
6, Absorption curves for cellophane and polythene••••• 21
7, Reduced pressure evaporation a9paratus ••••.••••.••• 23
8, Decay plot of hafnium isotope 181•.......•••••.•••• 24
9, Counter circuit With shielded semple container••••• 26
10, Crossection of shielded sample oontainer••••••••••• 27
11, Constant temperature bath••••••.• , •••••••••.••••••• 30
12. So~ubility plot of hafnium oxide in aqueous
solution••••••••••. ~2
v
LIST OF TABLES
Table Page
1. Solubilities of hafnium containing compounds ••••••• 6
2. Solubility versus the reciprocal of absolute
temperature •••••• ~3
3. Experimental results of solubility measurements ••• 36
4. Solubility Product of compounds at 25 deg. 0•.•••• 41
r
INTRODUCTION
To date, there is no information in the open liter
ature concerning the solubility of hafnium oxide in water
other than the fact that it is very insoluble. (1) This
(1) L nge,N.A., Handbook of Che istry, 6th Ed. Sandusky, O.andbook Publ., 1946 p. 200.
lack of experimental information is due, partly, to the
fact that hafnium is a relatively new addition to the list
of known elements. It was first isolated (2) in 1923 by
(2) Hevesy, G. and Coster, NatUl~e, Ill, 79.
Hevesey and Coster from a zirconium mineral. The extreme
difficulty of separating hafnium from zirconium has greatly
retarded research on the chemical and physical properties
of hafnium. Almost all zirconium minerals contain from 1 to
24 per cent (usually 5%) of hafnium. The early methods of
separation of hafnium and zirconium ere by the fractional
crystalizatlon of the double alkali fluorides, ferrocyanides
and oxalates (3) in which case the hafnium remained in the
(3) Ephraim, F., Inorganic Chemistry, 5th Ed., N.Y.C., Interscience Publ., 1948. p. 450.
mother liquor. The modern method is by ion exchange resins.
2
The mechanism by which a solid dissolves is, in its
self, an intrigUing process. A substance is said to be
insolRble because of the relatively lar~e amount of energy
required to separate individual molecules or atons from
the crystal form, or to ionize it. This energy is usually
supplied in the form of heat, hence, the solubility most
often increases TIith an increase in temperature. The re
quirement in energy may be classified into two increments;
first, that which is needed to separate the molecule from
the crystal, and secondly, that which divides the molecule
into ions. Therefore, the gereater the bond enerey between
ions of a molecule, or molecules of a crystRl, the less
tendency there is for the substw1ce to ~o into solution.
It is possible that the hafni~l oxide molecule ion
izes in the following manner:
Hafnium oxide is so !msolnble ~n aqueous solution that
conventional quantitative oethods for the detenaination of
its solubility cannot be employed. Prior to the introduction
of tracer chemistry, the solubility of many compounds was
qualitatively ~redicted from periodio properties. At present
the Atomic Energy Commission lUiS made available for experi
mental purposes, an artifioially produced radioactive
3
hafnium isotope with an atomic weight of 181. It was with
the oxide of this isotope that the following investigation
was made.
As no information exists concerning the exact solu
bility of the oxides of any of the members of Family A in
Group 4 aside from the fact that they are all very insol
uble"and the upper limit of the solubility of thorium-5
oxide (5) is less than 2 x 10 gm per liter, the solubility
of one of these elements would sive some basis for estim-
(5) Seidell, Solubilities of Inorg. & lfetal'Org. Conpounds,I 3rd. Ed. N.Y.C., Van rostrand 1946. p.1533
ating the solubility of the others. Aside from this benefit
tracer chemistry opens up a new micro-analysis technique
that exceeds in facility and accuracy the identification or
elements in extremely dilute solutions containing of the
order of a million atoms. But, being a new method, the des
cription of suitable apparatus is limited and it is hoped
that some of the arrangements used herein may be of future
use.
Recently, great interest has been evinced in the min
eral resources of the oceans of the world. Taken as a whole
they contain more precious metals, more alkali metals and
earths than have yet been extracted from the land. Those
present in abundance such as mngnesium are being recovered
on a large scale. Unfortunately, no prooess has yet been
deVised for the profitable extraction of those elements
present in minute quantities per unit volume. This ~ay be
due, at lEast partly, to the difficulty of laboratory an
alysis of the samples. Tracer chemistry is readily adapt
able to problems of this nature, since solutions of less
than one hundred-millionth molar are easily analized by
means of appropriate radioactive tracers.
4
5
REVIEW OF LITERATURE
A search of the literature failed to give any indica
tion of quantitative data. on the solubility of hafnium
oxide in aqueous solution. Furthermore, no exact informaA
tion was available on the oxides of titanium, zirconit~ or
thorium, the other members of Family A in Group 4. (6)
(6) Chemical Abstracts, 1927-1949.
The various Chemical Handbooks (7) (8) (9) merely list
(71 Lange, Ope oit. p.l
(8) Perry, J .H., Chern. Engg. Handbook, 2nd. Ed., II.Y. C. ,McGraw Hill, 1941. pp. 333-367 •.
(9) Hodgman, M.S., Handbook of Chern. & Phys., 31st. Ed.,Cleveland, Chern. Rubber Publ. Co., 1949. PP. 447-548.
hafnium oxide as very insoluble. Seidell (10) lists the
(10) Se~dell, A., Solubilities of Inorg. & Metal Org. Compounds, Vol. I, 3rd. Ed., N.Y.C., Van Nostrand, 1946PP. 604-5.
solubilities of other compounds of hafnium (see table 1)
It was further noted that the literature did not
yield any techniques for the handling of radioactive iso
topes in evaporation techniques. Therefore, those methods
tried by the author and discarded, may very well find ap
plioation when investigating the solubility of.a substanoe
haVing a greater solubility than hafnium oXide.
TABLE I
The Solubilities of Hafnium-contained Compounds in Acid
Solution
6
Compound Solvent Temp. Reported as Gins/liter
HfOBr2 13.36 N 25 C. Hf02/1iter 0.80 gInaHBr
II 8.77 N 11 " 10.6 GmeHBr
HfOF2 1.06 N 11 11 568.3 GmsHF
" 6.03 N " " 892.3 GmeHF
HfOC12 5.64 N 20 c. HfOClr:,fliter 0.16 GinsHCl
(~
HfK2F6 .125 N 11 HfK2F6/1iter 67.3 GrosHi'
HfO (H2P04) 5.94 N II HfO (H2P04) Iliter 0.026 GinsHCl
EQUIPMENT
COUNTING EQUIPMENT
A Berkley Decimal Counter, model 1000-B was used
for the first part of the investigation.
A Nuclee.r Corporation counting circuit, lllodel 161,
eqUipped with a Cenco register was used for later work~
The Counter tube used for the above two counting
circuits was a Victoreen Thyrode tUbe, model IB-85, op
erating at 900 volts. (see fig. 4)
A Lauritsen Electroscope, modell, eqUipped With a
charging circuit was used. A Plexiglass sample holder
was made and secured beneath the ionixation chamber of
the electroBcope.(see fig. 2)
A Palo-Meyers stop-watch was used to time runs.
This watch has a least count of 0.2 seconds.
EVAPORATION EQ,UIPMENT
One dozen 75 ml glass electrolytic beakers were used
to evaporate solutions saturated With hafnium oxide.
A 50 rol glass pipette was used to remove and trans
fer the saturated solution to the evaporation beakers.
A 2 liter Florence flask was used to make up and
store a mixture of visibly excess hafnium oxide in water
which was held at selected temperatures for equilibrium
to be reached.
Evaporations were carried out using an Acme Eleotrio
Co. hotplate, type 14221, operating on 110 ac.
The constant temperature bath was made trom an 18
7
8
liter glass jar, 12 inches in diruneter. This jar was placed
in a wooden box and insulated with six layers of asbestos
paper and a minimum of two inches of exfoliated vermiculite.
The jar was filled with water and eauiDPed with::l. _ -
1. A 0.1 deG~ee Gent. thermometer calibrated from zero
to 100 degrees Gent.
2. An ~ninco mercurial thermostat, model 930-07.
3. A Genco electronic relay, model 99782.
4. Three bayonette type heating elenents,Genco, 250 W
5. A Genco electric stirrer, type 50/60/.
A l/sth inch layer of mineral oil ~as poured on the
surface of the \'later in the bath to prevent evaporation at
the higher temperatures.
Weighings were made on a Seeder-Kohlbusch analyetioal
balance.
9
CRE aCALS USED
Hafniu.rn Oxide
A 0.9 gram sample of hafnium oXide, containing radio
active hafniUIYl was obtained from the Atom:l.c Energy Commis
sion at Oak Ridee, Tennessee. This sample was produced in
the Nuclear Reactor (see fig. 1) and when measured Just
prior to shipment, on June 16,1949, at 10 am it had an ao
tivity of approximately 50 me. This is equivelent to:
1.85 x 109 Beta emissions per seo.
From this data one may calculate that one atom of
hafnium in each 250, 000 was initially radioactive.
The half-life of hafnium 181 i8 46 days. It deoays
with a Beta energy of 0.64 mev and a Gamma energy of from
0.13 to 0.60 mev.(11) Hafnium 181 decays 99~99% into stable
(11) General Electric Researoh Laboratory, Chart of theIsotopes, 1948~
tantalum 181.
Hafnium 181 (uns~able) 46 days
~9.9c;;J%Beta
.64 mevTantalum 181 (stable)
.01%Beta
.52 mevTantalum 181
!22 us
ITe-.ll mevgamma .47 mev
TUngsten 181 (unstab~e)
! Kno betagrunna 1.8 mev
.-T..;..;un;.;aglll.,;s;;...t-.;e;.;.no.....;;;;l...8,;.-2 (s t able)
6e2
~ ~B~5P?5d2
4e14p~4d:;04rl4
+onJ.~
e2
~~5Sf5p~5d2
4p~4d:;04f14
+ 0
In nucleus: 108 n, 72 P
F1g. 1
In nucleus: 109 n, 72 P
The Formation of Unstable Ho.tn1wn 181 from a :Neutron-Gamma Reaction
J-lo
11
The results ot a spectrob~aphic an8~ysis of the hafnium
181 used in this experiment may be seen in Table I of the
Appendix.
Sodium Sulfate
Anhydrous sodium sulfate (C.P.) was used in the deter
mination of the unco~rr9n ion effect on the solubility of
hafnium oxide in water.
Ammonium Nitrate (a)
C.~. ~monium nitrate (a) which is stable from -16 to
32 deg. C. was also used in the determination of the un
common ion ef-fect on the solubility of hafnium oxide in water.
12
EXPERIMENTAL 1,':ORK
The purpose of the investigation was to determine
how much hafnium oXide dissolved in water at selected
temperatures. Of the various procedures possible, the
best appeared to be to evaporate a known volume of water
saturated Y]i th racl10acti ve hafnium oxide and then deter
mine the amount of radioactivity of the residue remain
ing in the container. Corrections for the natural radio
activity of the container and the background would be
required. The activity of a known weiGht of hafnium oxide
measured under identical conditions would then perrnit a
relation to be established between the weight of hafnium
oxide and activity measured.
Geiger counter tubes are known to become erratic,
and for this reason i tX!QS decided to use a Lauritsen
Electroscope (see fig. 2) to check the counter tube from
time to time.
The first task lay in calibrating the instnments
to be used. The sensitivity of the Lauritsen Electroscope
was plotted over the range of scale positions (see fig.3)
and from this ulot the optimum operating range was seen
to be from 30 to50 on the Beale. Next, the rate of discharge
of the electroscope due to cosmic rays and leakage of the
applied potential was measured. &ix, 30 minute runs were
made and an average background reading found to be 0.055
divisions per minute. Since the eleotroscope does not give
13
Fig. 2
Lauritsen Electroscope, Bureau of Standards Sample
ni Sample of Radioactlve Hafnium Oxide i th Lead
Container
14
-
I- --
- --+.-
90
i
L1----
80
I I I_-L1- _.l---i- l"t I r
--'- -
70
II
~-_·t
60
.•ti
,. i !~;-,-W:.:- +1-__ ~ -_,. .., rt -- - - I,. . - 1
50
. If:- i 1- •->---+Ir-.+----I--~,-+----...-11-
• ! t t
I'
40
-- -r-r
-
30
•I
!-
20;
--i----
.
-- I-
-I-
,
-
-Tr-j': Li~es
i
-
I·
--l-~---t -10
..--r-
I
1--- f- -
- -..--1-
--
--- .--J~-1t_--'-_I---_.,......-1----:I~+--
1
-
o
60
50 I,
90
30'
20
10
70
80
130
140 t~_.....---,..I-..------,--y----r--:--~_:______,-...,____r___:_-..______._---.or~-:---r--:-.~,t-_':"T+I-:-:t-.7:'"r:::;"'~4'~.t-:-_-'7"""~~"'7:-'T"':'H:-i-:r--,:---.-----rL---,r-.,--.......-r.~~r---.-...,.-r-II'-=--or-,
-1·-~-1~-_~-+---+-I~'.j--+---I,-+-+--.J!=---r--!-"'--l---+-:~"""+--l!-·-·+-"-·+--+--t--l---.... -~--1--+-__ +-'-+-1- -1--+--+-,r.IL . • t. T I
t 1- I I- -1'-+--+---+-+-+--+1-1-:-1 ..--+-+-+--+-~-1----+---I-+--1-..--+!~t- -:-,r-l-I,---t-.-+---::....--l--.l-
I
-i'---;--t--+-+--r--t-t--+---+-~i---+--j
- -I--~ I.. Ii: +----4--+-.-11-:+~~'-...:..,'--I--I--;-I1-+---I- - - - -+---+--..l...--1-+--+--~i--i---l-I-l
..-- - -- -I----+-'-+--;'--+-i---+---i---I-",:,,--+--;.I
-1120 [-
100
ttlrd 110~ooQ)
tr.l
Scale Positions on Electroscope
Fig. 3 Plot Showing the SensitiVity as a Function
of Scale Position
15
direct activity measurements (but only the correlation
betveen intensity of activity and collapse rate of the
quartz fiber) it vas standardized wit~ a Bureau of StanA
dards smnple of natur~l radium D Dnd E clectroulated on
a palladiuJ11 coated 8ilvel' disc. With the e:;Lectroscope
thus calibrated, a rapid check may be kept on the counter
tube. The Lauritsen Electroscope, while reliable, is not
sensitive enough to give the results obtainable with the
counter cirelli t, hence was only us ed as a checle.
The Geiger counter circuit was also calibrated. The
09timum operating potential of the Thyrode tube was found
by plotting the activi ty of D. racl10active sample versus the
voltage applied to the tube. At one region of the curve,
(875 to 900 volts) a plateau was reached (see fig. 4). ThiS
is the optimum operating voltage because over the plateau
range the number of counts per unit time is independent of
small voltage variations. Settings may be made anywhere
within this range for reproduc~ble results. A check made
on the background actiVity of the tube Without an' shield
ing yielded an average of from 57 to 63 counts per minute,
al though this figu.re may vary daily and even hourly depend
ing for example, on climatic conditions.
For the first experimental run, a tube holder was
deVised so that the counter tube was held in a horizontal
position (see fig. 5) 6 em from the sample position. A
standard sample containing 0.00030 gms of radioactive oxide
was nlaced on Et 1 inch watch-glass and the actiVity measured•..A small runoun't 'of the radioactive. oXllde in visible excess was
f-- I
------t-- -~-
! I! I , =r-'~
I- --- .I
~. .. E
-I •. -. -- --H ' 1-- - . ~-1 =--~--=-=- -~:: If - . =- )
I I +--'- ~ - - t=-: -~- ~~,~~ - : ;~' p-~~~eau- ~~1-Nl ' ' !---,- - ".'1800 I I .. I I
1900
--
~ 4'-$'~-- ,--.. :: --.~. --~.. - - . - ,-._ r --
- ~-! -- 1--- -, ~-. ---
1---- --, I
1--..-- 1·----
~.-~-~~,'r=-~: --"~- , -,.__ - - - - . _' - - .~ - - I _ .....
---- --- - -'--- --1700
f--~- -
- -+------ - -.. I
~ - _. ·_t . _
_. - ... - ----,- t~-E--'-,- -.- .,., .._ .... _~_ .. ~ - ........ 0 __-
'-_. -'--_. ----" ~~ --:- --- _..-----+------- --.
"-H j-'~-/=r:. -'..-_ .._--
t~ /'
~--:--r: I ! I
I-- I
'-~ II -1. , ' :' - - --; --I-- ~ ~~._,_,~
. ,-._- ---=t=--~ 1= ---"'--1---1.----- 1 - -- --I I : . . j I I... ~ ~ ~ I -t... ... ..:.. __'__. _ '!"" .. .. _
- - --- - ---il~~•• _ : I ' • -=,"_-=-.-= --;-{- - e-yon --950~-t,S. -i.s.
1=----. - - -,. -:- .' .;.-~.-:.- ' :<, .• r1:tlNIl J' gloll...o:r.:·__ I I ~~,- - .- ~--:-....- - - -tub ---
~-,- : .._! __._;_,_-_- 1 .. • ~ -= _-_-.:- _.
I -- - ..-
. .. , ,-- ...--
Q).p~s::~
~ 1600HQ)
PI
Ol.ps::~
g 1500
I; I I I I iii I I I ! I ' I ! ' I I" I I I' '~I" r .,.....- - - I-+-i • [i i j. i '. ----- ~ .. -- - -
400 I I I I I I I I I ' . , 1 ' , ' 'u " u , I ' - i - '/- 1-' I, " ,1 . I i I I I 'I I Iii, . , i I Iii i I ! t, i -
Applied Voltage for Victoreen Thyrode Tube, Using a Nuclear
Corporation Count1ng C1rcuit.
......())
950900850Voltage Applied to Tube
Plot ShoWing Change if Counts per I.:inute With Increase of
800
Fig. 4
To lu.~h volta~esupply
COlmter tube
Fig. 5
Tube mount
Sample position
17
Horizontal Counter Tube Support
18
placed in an 8 inch test tube containing 25 ml of distilled
"lU tel'. The test tube was held at 26;t .3 deg. C. for one
week with occasional agitation. At the end of the week, a
2 ml portion of the supernatant liquid was pipet ted from
the test tube to a 1 inch watch glass, 0.5 ml at a time,
and evaporated to dryness on an electric hotplate. The ac
tiVity of the residue was not statistically significant
when the background was subtracted. This indicated that the
solubility of hafnium oxide was so low that an insufficient
volume of saturated solution had been evaporated. For a sam
ple to be judged to have a significant activity, the activity
must be greater than the backGround activity uncertainty
which istiN, where N is the total counts per time measured
for the background activity. For example, if the background
activity of a counting circuit were 60 counts in a centain
time interVal; ant sample showing a net activity of 7.0
counts in the same time interval could not be considered as
being radioactive sinoeti6Q ist??
The volume of solution evaporated in the above manner
was increased to 4 ml and finally to 6 ml without yielding
ant valid net activity. From these runs it was apparent that
the extreme insolubility of hafnium oxide demanded that
larger volumes of solution be used; therefore this method
was abandone.d.
The seoond experiment was set up to accommodate 50 ml
of hafnium oxide saturated water. The solution to be evap
orated was allowed to drop from a burette onto a heated 2
inch watch glass. It took 36 hours to evaporate 50 ml of
19
solution in this manner. Three separate samples were taken
at 30 deg. C. and after evaporating t 1em to dl"'yness, they
were nlaced under the counting tube. No statistically sig
nificant activity above the background count was noted.
Since this apparatus could not be enlarged without involv
ing runs taking several days, lttoo was 'abandoned and an
improved method sought.
From the data obtained it was see that a method must
be devised to fulfill the following requirements:
(1) Allow an evaporation of 250 ml or more to a
small (1 inch dia.) v!orking surface.
(2) !lake posGible an ovornight evaporation proceS8.
(3) Permit an increase in solubili t Jr by physical
m ana (terperature increase).
(4) Allow no decrease in accuracy.
An attempt to satisfy these conditions was made in
the folIo "ing m nner. About 500 ml of distilled water was
placed in ~ one liter Florence flask and to this was added
a visible excess of radioactive hafniwn oXide. The mixture
was boiled and allowed to cool, then placed on a hotplate
set for 60 deg. C. and kent for ten hours. The stem was
cut from a four inch glass funnel and the apex sealed with
a flame. A cellophane cone was folded and olaced in the
funnel cone and to this was added 5 suoce8~ive 50 ml port
ions of the previously made solution, while heating the
funnel and contents on a sand bath. It was found, however,
that above 75 deg. C. the cellOphane became porous, allow
ing a seepage of solution. In an attempt to locate Borne
cellophane that would remain waterproof, a letter was
sent to the Cellophane Division of the Du Pont Co. In
reply it was stated that no water)roof cellophane was
avo.ilable, but it was suggested that I substitute Poly-·
thene for my work. When this was used in nlace of cel
lo.hane, and 250 ml of hafnium oxide contained solution
evaporated, no seepage occurred. It was hoped that the
Polythene could be folded and then dissolved upon a
watchglass and ignited to an ash, leaving the evaporat
ion residue V11ich could have bean tested for radio
activity. However, when the Polythene cone was folded
and pla.ced on the watchglass encl the toluene was adcled,
a gel formed. Upon ignition of the gel, violent sputter
ing occurred and much of the sample was lost.
It i8 believed that of the two, the cellouhane
would have been the better except for its porosity. Ex
perience shovled that upon completion of the evaporation,
the cellophane could easily be dissolved with acetone.
No undesirable gel fommed and ignition was easily con
trolled preventing loss of the sample.
To ascertain which of the two acted as the greater
shield to Beta radiation of the hafnium oxide, tests ~ere
run using varying thicknesses of Polythene and cellophane
(see fig. 6). It was found that Polythene had the least
shielding effect on the basis of mg/cm2.
Although the technique of evaporating a large vol
ume of saturated hafnium oxide solution in an envelope
which could be removed and dissolved With the proper
20
250 -=::1'" g- If ,,:, f---=r==--~--I j - • i I I. t I ----1--
-----.- -. =t-- .--,--- -
~ -
w:- :mi~. .,~'- " . ~ I III-' -, <>~J ~~-:~:::-j-'---.~--
.......
II!!IIII I-~' I".~
-=±:e±h I: : :~ ~--
-- I
= =:L-=- ==--.-1------
--l-.-LI I I I F I I " ! F ~~_,_. _., ,I.. ~
, I I I, I
I- ~ ... - ... - -
- I ':':~~~~~-F-; ~:~ __1· . I , , •I
I I I I
': . I I , - ,- --- - - I,. . .~ -. --- --t--- I . I - - --~. ,_ I ' "H" - - I - - .--h-i-;~ I II i±iiE I -;,~=:: ~~ :~':.1 _:-:- :.. ! II' . I • . - ." ,I - --- - .. ... +- -: >I .-.- -=-- -''-- -
50
5 10 15 00 25 30mg/om2 ot absorber
~he Absorption of Hafnium 181 Radiation by Different
Thioknesses ot Cellophane and Polythene
oo
I
Fig. 6
~TT ,-,-
- ~ I ' ... I
35
r-uI-'
22
solvent, and subsequently ashed and then tested for radio
activity did not prove sati'factory experimentally, it was
not abandoned. If reduced pressure Were employed to in
crease the rate of evaporation of solution, the cellophane
might still be used. The srumple was placed in a glass des
sicator (see fig. 7) which was connected to an aspirator.
The reduction of yressure caused the solution to boil
violently, forming bubbles the full diameter of the funnel
and much of the solution ";as lost. \'fhen the nressure within
the dessicator was increased slightly, the boiltng ceased
and only 25 ml l'/ere evaporated in a 10 hour run. Uethyl
alcohol was a(1(1ed. to the sa.rr1"ole to decrease the boiling
point, but this 11a<1 no apparent effect on the solution's
evaporation rate. The dessicator was heated on a sand bath
and reduced pressure evaDoration continued, but the dea
sicator cracked. Since entirely satisfactory equipment for
a reduced 9ressure evaporation was unavailable, the pro
cess was abandoned.
Keeping in mind the previously stated requirements for
a practical evaporation techniqlle, the final process was
eVolved and used in the actual solubility D.Ci38Urements. It
is important to note that by this time the strength of the
original activated hafnium oxide sample had decreased con
siderablY (see f.ig. 8) so it was considered nece8s~y to cut
dovm the background activi ty count and to bring the counter
tube closer to the sample position. To satisfy these condit
ions, a combination sample holder and lead shield was con
structed. This arranGement nerm1ts the duplication of
Base otdessloator
aspirator
23
Fig. 7
Apparatus tor Evaporat1ng 250 ml Portions ot Saturated
Solutions of Ht0Z in Oellophane or Polpthene Envelopes
24
.
.. 1---+--1---t--I--t--I- -.--
..- - - - -0-~ - _ -- r--- .
I
1:.1 I---~-·--·I- 1- . - I--'~
10 ,.
I! I ~ !.!-I-9I'~; I .- -c-'-+-+j-+--+--I---+--l
• : I I() ,1 I I . f- - 0-
~ 8 t - 0, '~--t---t--+--+--t---+-I--+--+--+---I
~ 7 ! I _ ~! ~~ G - -- - --~ I I HI . --- -""r---r-.- I~. 6 '-----+-+-+--+---+1
1_ _ _ r--k:>
~ 5 L..---+---+_--t--0
_j_ i I !... I-, -c'. . - c- ....... r'- ~I ---r-~ I Iii~ 4 ---j-i--I I [--j - -- -- -~ 1--+--+--7'i--+---+-' - - -i--I-------+-+---!---I----f--j---l--+--+-+--1--l~. II~
:>::13~IHo 2
o
!', , I. I_."';:- •....,.--j.-t-'--+--- --1---1---1----1-- ,__oj,-- i,- I
I' :i'. t.I ~ ~ l-rl' : 1 .:; • ~
o 100 200 300 400 50u
Fig. 8
time in days
A Plot Showing the Logarithm of Activity
in Beta particles/ second as a Function
of Time for the 0.9 gm Sample of Radio-
active Hafniu~ OXide
25
posi tion of counter tube and sample (see figs. 9&10). The
geometry of position is important v!hen making a comparison
of sC'J11ples, as the change in projection of activi ty on the
tube would result in a change in recorded activity.
A sct of electrolytic beakers were obtained, so
chosen because their inside diameter was just 0.5 em larger
than the counter tube, and their length about the same as
that of the tube. A lead shield was cast to permit these
beakers to be slipped inside and provide one inch of lead
shielding from outside radiation. A lead plate 0.5 inch
thick was drilled to permit passage of tube and holder, yet
removable, so that the beakers could be easily extracted.
The plate to which the counter tube is connected was con
structed of' 0.25 inch aluminum which adds to the shielding
effect. When evaporation is carried out in the beakers, the
residue will either cling to the inner Walls of the glasB,
or deposite on the bottom; in either event, this residue
containing the radioactive hafnium oxi~e will be close to
the tube and any radiation from it will be readily recorded
by the tube With a mimimum loss from random scattering.
The background actiVity in the above described count
er shield would be a combination of whatever outside acti
Vity got through the lead barrier pl~s any actiVity in the
glass beakers. This compound background was measured and
recorded for each of' the empty and cleaned beakers and it
varied from 58 to 75 counts per minute. Sinoe the baokground
Without the beakers was 40 oounts per minute, the added
actiVity was attributed to the presenoe of a radioaotive
Fig. 9
Counting Cirouit and Sample Holder
Showing the author inserting an evaporation bewcer into the sample
oontainer, prior to oheoking its aotivity.
~~
27
Counter tube
Evapore.t1onbeaker
Lead shield
F1g. 10
Cross-seotion diagram ot oo~btnatlon lead shield and sample
holder.
28
element in the constituents of the glass. Several other
pieces 0 glassware were cheelced for activi ty a.nd it vias
found that all those tested showed some activity, nlthough
pyrex glass seemed to contain the least. Since all beakers
and glassware were cleaned With hot cleaning solution
( H2S04 & K2Cr20?) it is unlikely that this activity ~as
from surface film.
'When it was decided that 500 ml of solution would be
evaporated in ten sjlcceasive 50 ml portions, a 2 liter
Florence flask was filled with distilled water and a visible
exoess of radiated hafnium oxide added. The mixture was
stirred well and the open neck covered With a 50 ml ~yrex
beaker, effectively keeping out duct but allowing for ex
pansion of air within.
A constant temperature bath was constructed as des-
cribed under Equipment, on page 8. The 2 liter Florence
flask containing the hafnium oxide and water mixture was
immersed in the bath at a selected temoerature for a oer-~ -
iod of 24 hours to insure maximum saturation. It was
stirred frequently, the last stirring being four hours
prior to taking the first portion, thereby allowing the
settling of any partioles of hafnium oxide not actually
in solution.
~hile the solution was reaching equilibrium, a
standard sample v.as prepared in the following manner. A
small amount of hafnium oxide (0.00026 grns) was weighed
out and placed in one of the evaporating beakers marked
29
II standard". Extreme care must be exercised in handling
and transferring this sample because the accuracy of the
entire experiment hinges on the dependability of the
standard sample. To this beaker containing the hafnium.
oxide were added several drops of distilled water and I ml
of concentrated nitric acid.
Hf02 ... 4HN03 -. Hf(N03)4+ 4H20
The volume was brou ht up to 50 ml with distilled
\'later to duplicate the volume of the samples to be taken,
The solution WEtS then evan orated to dryness on a hotplate
and fitted with an aluminum foil cap to protect it from
the surroundings. To ascertain the actiVity of this stand
ard sample, the cap was removed and the beaker placed in
the counting container. mhe plate was placed over it and
the tube lowered into position. A total actiVity of 287!3
counts per minute was recorded, and the background ot 6l t J
counts per minute being subtracted Ie t a net actiVity of
226 counts per minute. This value will denreciate as the
hafnium decays.
After the 24 hour waiting period, the solution in
the bath (see fig. 11) vIas assumed to have reached equi
librium and, noting time and temperature, the first 50 ml
portion was taken by means of the filter eqUipped 50 ml
pipette. This portion was rapidly transferred to one of
the evaporating beakers which was then placed on the
hotplate for evaporation. The ptocedure was duplicated for
each of the other 2 beakers (three samples per run).
Fig. 11
Constant Temperature Bath and Evaporation Equip
ment Including Hotplate, 2 Liter Florence Flask
and 50 ml Pipette 1th Attached Cotton Filter
31
Before the beakers had been completely evanorated to
dryness, successive portions were added to prevent Bud
den cooling and consequent cracking of the beakers.
~ith the taking of each set of portions, the time and
temperature vrere recorded so that an average temperature
could be arrived at; there being a 0.3 deg. C. deviation
in the thermostat.
Upon completion of the 500 ml evaporation of the 3
samples, they were allowed to cool, then separately placed
in the countine shield. When their respective actiVities
had been counted over a 30 minute period, the background
was subtracted from each and this net activity compared
with that of the standard sample. The solubility at this
particular temperature was found by means of the follow
ing relation:
weifht of HfO~ in Stanclard Sample _ WCi~ht of Hf02 ResiduectivIty 0 Standard Sample - Act vity of ResIdue
The actiVities found for each of the three samples
were averaged and the logarithm of the solubility found
was plotted (see fig. 12) against the reciprocal of the
absolute temperature of the run. Using the above tech
nique, runs were made at 35, 50, 60, 70, 80, 90 and 98
deg. C. (see table II). Taking the points in figure 12
as a straight line function, a line was drawn most
closely representing the solubilities listed.
Since there were only a dozen evaporating bew{crs
available, it was necessary to use each several times. To
f ] I
I
2.7 2.9 3.1lIT xlO 3 deg. Kelvin
Fig. 12 Plot of Logarithm of Solubility of Hafnium
32
-7.00 -r I - -I -T 1( !II I
-6.95 --r -!-+++--ej
-6. 60 _..!--....L--L...--I--.!.---L--I.---I_l...-J..-.L-...L--!--L.---.!---l.---L---l..._-->
2.5
-6.90 ,
§'2 -6.75
~J I..c:0)
~ Iof3 -6.70·
Cl-Io
UJ:3 I __
-6.65 --J
ooo<-6.80Q)
<d.,..;>-~o
Oxide Versus The Reciprocal of Absolute
Temperature
TABLE II
Data For Figure 12
Solubility in Log of Absolute Reciprooal ofMoles/lOOO f::,ms Solubility Temperature Absolute Temp.
Water deg. K
1~ 07xlO-7 -6.9706 307.9 3. 24xlO-3
1~33 II -6.8762 323.0 3.10 II
1,36 " -6,8665 333,3 3,00 II
1.51 " -6.8211 346,6 2.88 •*8,31xlO-8
-7.0804 353,6 2.83 "-7 -6,6596 363,0 2.75 n2,19xl0
1.75 " -6.7570 371.3 2.69 II
* ThiS data not plotted.
33
34
prevent an accumulation of activity, the beakers were cleaned
with hot cleaning solution before being used again. When
tested, it was found that a be~{er might be 5 per cent more
radioactiVe than when in1tially examined but it \vas believed
that the activity gained was well imbedded in the be~cer and
corrections were made for this activity.
The results of the experiment carried out are record
ed in Table III and graphed in figure 12.
It was decided to add uncommon ions to the hafnium
oxide solution at about 60 deg. C. Into a 500 ml volumetric
flask Vias placed a Visible excess of hafnium oxide, and 200
m1 of distilled water were added. Several grams of C.P. an
hydrous sodium sulfate were placed in a clean 250 ml pyrex
bea~er.and heated for an hour on a hot9late at about 150
deg. C. to drive off any accumulated water of crystalliza
tion. After allOWing the 80cliwn sulfate to cool, 0.7100 gm
was weighed out on the analytical balence and carefully
transferred to the volumetric flask cont~ininG the hafnium
oxide and water. After dissolVing the sulfate in the 500 ml
flask, the volume was brought up to 500 ml and the flask
was immersed in the constant temperature bath, which was set
for 64.4 deg. C. The solution was given occasional stirring
and allow;bd 24 hours to come to eqmlibrium. A clean evap
orating beaker was calibrated for baokground actiVity and
placed on the hotplate. Four, 50 ml portions of the equi
librium solution were transferred to the beaker which was
subsequently evaporated to dryness using very little heat to
35
prevent splattering. ~~en checked, the beaker containing the
sodium sulfate and residue showed an increase in hafnium
oxide over the amount indicated by figure 12 for 64.4 deg.
c.A 0.1 Molar solution of C.P. ammonium nitrate (a) was
made up in the same manner and the erperimentally determined
inore~se in hafnium oxide present compared with figure 12 as
above.Data recorded for this and the sodlun sulfate add~tion
to the hafnium oxide-water mixture appears in ~ab1e III.
TABLE III
EA'"Perimental Results of Solubility Measurements
No. Temp. of Time jor Net cts Av. cts Vol of Conversion fector Ion. Solubility inHf02 soln sat 11'1 in per min per min sol'n fll 6 fl2 H3 Stl'. gIns of Hf02t'lOOdeg. C. hours residue. residue taken x10- .... ,'" gIns water. ..._~
ml
1 26 .3 168 0 2
2 It II 1 2 0.5 2 2
3 n n 2 2 4 Solubility not calculated since
2activity found was not statis-
4 II II 0 1.0 4 cally significant.
5 II II 0 6,
6 II n 1 2 0.5 2 6
7 30 2 10 sample lost 50
8 60 2 10 sample lost 50
9 60 2 10 sample lost 50
10 60 2 10 samtlle lost 250,
11 70.3 ,3 24 17.3 It I..
12 II u 12.7 li 101 14.0 It 500 1.10 .20 .9~7 ;3,~23xlO-SI
13 II II 1~.0 Ii ~
14 80.3 .3 II 6.3 1-?J
~ 7.1 1-1l15 1/ II 7.8 Ii 500 1.16 .20 .971 1~77xlO-6
16 n 11 o Ii17 89.7 .3 II 17.0 l·~
18 " If 19,4 It ~18.2 It 500 1.22 .20 .965 4. 67xl0-6
19 98.0 .3 II 13.4 l·~ I
20 " II 13.4 It 13.5 It 500 1.32 .20 .959 3. 74xlO-6
21 " " 13~9 It22 49.7 .3 • 11 6.8 It -
23 ft II 5,4 l-l " 6.1 Ii 300 1.83 .33 .988 2. 84xlO-6
24 60~0 .3 II 4.2 It )25 II II 4.0 -It t 4.1 1~ 500 1.41 .20 .983 2. 91xlO-6
26 34.6 .3 n 8.3 l~' )
27 u II 7.1vl.;!s .. 7 7 1-~~ 500 1.48 .20 .997 2. 28xlO-6'" • 2
64.4 .3* II 6.4 It 6 4 11-· 1.94 .980...6
28• 2
200 .50 .3 6.60y..10
29 II ** II 4.5 It 4.5 It 200 f.OO .50 .980 .1 3.08xlO-6
* run made using .10 molar sodium sulfate
** a " " II II ammonium nitrate
CAen
DISCUSSION
(A) POSwibility of andom Particles
The relative agreement of the three samples Dcl~en at
each temperature run discounts the possibility of the ac
tiVity havin~ come from random particles sucked up by the
pipette and recorded in the reSidue. Had randomness been
the source of the activity, in all probability the plot of
the logarithm of solubility versus the reciprocal of the
absolute temperature would have consisted of isolated
points havtng no particular trend. Furthermore, there
would have been no reason for the solubility to show a
steady increase With increasing temperature.
(B) Possibility of Colloid :ormation
Concerning the possibility of the hafnium oxide hav
ing formed a colloidal suspension, Yeiser (14) states that
(14) Weiser, H.B., Colloid Chemistry, N.Y.C., Wiley, 1948.pp. 221-266.
when heat is applied to a sol, it tends to coagulate out.
Re further states that radiations such as x-rays, alpha
particles and ultra violet light have a sensitizing ef
fect on most sols. Furthermore, he says that the adQition
of an electrolyte has a coagulating effect on 80ls. It
seems improbable that the hafnium oxide added to distilled
water resulted in a colloidal suspension because the part
icles were visible andjlO- mechanism as present to red.uce
these particles to colloidal dimensions.
3?
38
(0) EVidence of a True Solution
Examination of the solubility-temperature data
(see Table II) indicates a fairly steady increase ot sol-
ubility with increase in temperature as 1s the case for
most solutions. Prom the data on the addition of an un-
common ion to the hafnium oxide-water mixture (see ~able
III) it is seen that the addition of an electrolyte in
creased the solUbility of the hafniun oxide when both the
sodium sulfate and the ammonitoo nitrate were added. As may
be seen in (E), the experimental increase in solubility
compared favorably With that predicted by the Debye-Huckel
'I'heory.
CD) Energy of Solution
The energy of solution (or heat of solution) of a
compound may be calculated from information on its solubil
ities at different temperatures from the follOWing
equation:
(1) LOg!2 =: 4H~T2-T, ~81 R T2 Tl 2.30
where:
S is the solubility in moles/lOOO gros H2O
T is the absolute temperature corresponding
to the solubility
~ is the energy of solution in calories/mole
R is a constant 1.987 cal/deg./mole
By taking values for the solubility and corresponding temp-
eratures from Table II, an_average value forAH was found
39
to be 3100 calories per mole. Another means of calculating
the energy of solution is shown by the relation:
(1) Log S = (H!2.,3R)(1/T) + C
,where:
S is the solubility in moles/lOOO b~S H2O
T is the corresponding temperature (absolute)
AlI is the energy of solution
R is a constant to 1.987 cal/deg/mole
C is a constant
In figure 12, the logarithm of the solubility is plotted
against the reciprocal of the absolute temperature. The
slOpe ot the line best fitting the points is 643 and the
energy of solution from this slope is 2958 cal/mole.
(E) Application of the Debye-Huckel Theory
From the Debye-Huckel theory one would predict that by
increasing the ionic strength of an aqueous solution in
equilibriillu with hafnium oxide, one would increase the sol
ubility of hafnium oxide. Experiments (see PP. 34&35) using
a one-two salt (Na2S04) and a one-one type salt (}lli4N03) to
increase the ionic strength of the aqueous solution in con
tact ith hafnium oxide demonstrated that the addition of
uncommon ions did increase the solubility of hafnium oxide,
and approximately as quantitatively theorized.
If we assume that solution occurs by the reaction:
40
~ -o-Rf=: 0 + II-OR ...,.Rf.. 0 .,. 20H
we may calculate the equilibrium ~o~8tant for this reaction
from the data collected at 64.4 deg. C. at which temperature
the effect of the ionic strength was investigated.
where K is the equilibrium constant and the other symbols
are the activities of the ionic or molecular species indi
cated, Since the experiments were carried out a.t atmospheric
pressure w1d the hafnium oxide is a solid, we may take the
actiVity of water and hafnium oxide to be unity.
Substituting the products of concentrations, CX1 (whioh were
experimentally measured) and the mean actiVity coefficient,
f, for the actiVities, we obtain:
(3)
when pure water was used to dissolve hafnium oxide, the
total concentration of ions such as HfO··, OH- and R~ was7 -7 -7
80 small (about 1.5xlO- I 4.7xlO and 1.7xlO molar
respectively) that the mean actiVity coefficient, f, of
hafnium oxide was 0.998. Then t was calculated by means ot
the Debye-Huckel equation:
(4).J..
Loe::- f. -A· Z ·z·u2c .. _
41
where A is a constant equal to 0.539 at 64.4 deg. C., z. is
the magnitude of the ch.r e on the hafnyl ion, z_ is the
magnitude of the charge on the hydroxyl ions and u is the
ionic strength equal to 2~Ci'ZI in which Ci is the concen
tration of each ion and Zi is the valence of each ion.
If we take f as unity, the equilibrium constant at 64.4
deg. C. is:
( ) 2 ( -7) ( -7) 2K:: CHf04otCOH- ~ 1.5xlO 4.7xlO-20
;: 3.3 x 10
This constant may be calculated for 25 deg. C. and compared
with the solubility product for other substanceS at 25 deg.
centigrade:
TABLE IV
Solubility Product of Various Substances at 25 deg. C.
Substance Solubility Product
AgOl 1. 56x10-10
-15Al(OH)3 3.7 xlO
HgBr2 8.0 xlO-20
-20*Hf02 2.6 xlO
Hglz 3.2 xlO-29
-49Ag2S 1.6 xlO
* Extrapolated.
42
If we consider the experiment at 64.4 deg. e., in
which 0.10 molar of sodium sulfate was substituted for pure
water, we have an ionic strength of 0.30 instead of 6.2xlO-7
as in pure water, and the mean activity coefficient, f, is
given by the Debye-Huckel formula:
(1)
( 2) or f =0.26
Using this value of f in equation (3) page 40, the solubility
of hafnium oxide predicted by theory is 7.09xIO-7 moles/liter
or 6.95xlO-? moles/loaO gros water. The experimentally obser
ved value, uncorrected for absorption of activity by the sod
ium sulfate, is 3.1Xla-? moles/IOOO gIDs water which is more
than twiae as large as for pure water but only about half as
great as theoretically predicted. This is in satisfactory
semi-quantitative agTeement With theory since correcting for
absorption of activity by the sodium sulfate would increase
the experimentally determined solubility.
To avoid the correotion of absorption of activity by
the salt used to supply ionic strength, ~on1um nitrate was
used. Before the residue was measured for activity, the beak
er containing the residual hafnium oxide and ammonium nitrate
was heated until the ammonium nitrate was driven off. In
this experiment, the conoentration of ammonium nitrate ~as
-70.10 molar, the Debye-Huckel theory solubility VTas 3.8xlO
-7molal and the experimentally determined solubility 2.lxlO
molal. The agreement between experim8nt and theory is better
43
in this experiment than in the sodium sulfate experiment.
The theory is known to hold best for ionic strengths of
less than 0.09, so if more time had been available,
exuerimente would have been carried out at lower ionic
strengths.
SUMMARY
(1) The solubility of hafnium oxide in aqueous solution
from 35 to 98 deg. C. was experimentally determined.
(2) The Solubility Conste.nt -:for hafnium oxide at 25 deg.
C. was determined and compared with that 01' other
compounds at the saoe temperature.
(3) The Heat of Solution of hafnium oxide was determined.
(4) The consideration of hafnium oxide forming a colloidal
suspension in water was discussed.
(5) The uncommon ion effedt on the solubility of hafnium
oxide was experimentally examined in a preliminary
manner.
44
45
SUGGESTIONS FOR FURTHER S UDY
Experiments such as Were carried out for the inves
tigation of the solubility of hafnium oxide could very well
be made for some of the other members of Family A in Group 4.
The Atomic Energy Cor.~ission at Oak Ridge, Tennessee,
lists for sale the following radioactive compounds:
Titanium Oxide:
Titanium oxide containing artifioiall produced titanium
51 having a half-life of 72 days, TIith a beta energy of 0.36
mev and a gamma energy of 1.0 meV is available. The relatively
long half-live of this isotope makes it a ood tracer for ex
perimentation. Upon decay, titanitUn 51 goes 10 stable vanadium
51 with no side deoay products, hence there is no danger of
conflicting radiation from the daughter.
Zirconium Oxide:
Although only the hydrOXide of zirc~n1um 95 is av il
able for sale; this, however could be converted to the oxide
by heating. Zirconium 95 has a relatively long half-life
also; it is 65 days. Upon decay, zirconium 95 goes to colum
bium 95 which is also radioactive, ving isotopes ith half
lives ot 90 hours and 35 days respectively. This fact, plus
the beta radiation of 0.39 mev (a relatively weak actiVity),
makes zirconium 95 a more difficult isotope to measure with
out the addition of radiation filters, than either hafnium
181 or titanium 51. Solubility measurements With zircon
ium 95 are by no means impossible if proper consideration
is made for the heterogeneous activity emitted.
The Atomic Energy Commission does not list for sale
any compounds of thorium, since this element has six nat
urally occurring radioactive isotopes. It is very diffi
emIt to seuarate isotopes of the same element.
Studies of the effect of ionic strength on the sol
ubilities of hafnium and titanium salts would also be of
interest.
46
47
APPENDIX
Impurities in the hafnium oxide sample used, accord
ing to spectrographic ane~ysis by Oak Ridge National Lab
oritory:
Impurlt~ Amount present
Zr02 0.75 %as Zr02
81 very weak
Cu trace
Ca faint trace
Fe II II
}. g If It
Zn very faint trace
Pb n 11 It
N1 II " If
Na It • •B n " ..
Ag It II II
Al
The follow1ng elements "'Jere sought, but not found:
Al B1 Mn 1
As Cd Mo V
Au Co Sb
Ba Cr Sn
Be L1 Ta
48
BIBLIOGRAPHY
(1) Atomic Energy Commission, Catalogue of Isotopes, OakRidge, Tennessee. pp. 1-32.
(2) Cork,J.M., Juc1ear Physics, N.Y.C. Van Nostrand, 1947pp. 116-138.
(3) Friedlander and Kennedy, Introduction to Radiochemistry,N.Y.C. Wiley, 1949. pp. 1-259. .
(4) Garner, C.S., Lauritsen Quartz- iber Electroscope,J.Chern. Ed. Vol. 26 pp. (1949)
(5) General Electric Research Laboratory, Chart 0 theIsotopes, 1948.
(6) G1asstone, D., Textbook of'PhYSical Chemistry, 2nd. d.N.Y.C., Van Nostrand, 1946. pp.l18-137•.
(7) Hodgman, .S., Handbook of Chem. and Phys., 31st. Ed.Cleveland, 0., Chem. Rubber Co. publ. 1949.pp. 447-548•.
(8) Lange, N.A., Handbook of Chemistry, 6th Ed. Sandusky, O.Handbook Publ. 1946. pp. 200-279.
(9) andeville, Scherb and Keighton, Radiations From Hf 181,Phys. Rev. Vol. 75. ~o. 2, 1949. p.221.
(10) Perry, J.R., Chemical Engg. Handbook, 2nd. Ed•.•Y.C.I~cGraw Hill, 1941. pp. 333-367.
(11) Prutton a~d ,arron, Fundamental Principles of PhysicalChe~istry, N.Y.C. Macmillan, 1944. pp. 465-504.
(12) Seidell, ~., So~ubilities of Inorg. and lietal Org. Compounds, Vol. I, 3rd. Ed., N.Y.C., Van Nostrand, 1940.pp/ 604-5.
(13) U.S. Commerce Dept., Safe Handling of Radioactive Isotopes, Handbook No. 42. 1947. 28 pp •.
(14) Weiser, H.B., Colloid Chemistry, 1st. Ed. .Y.a., iley,1948 pp. 72-83.
(15) Young, R.A., Vapor Pressure of Thorium Acety1acetonate,J. Am. Ihem. Soc., Vol. 61, p (1939)
Hampden O. Banks, Jr. was born November 12, 1922 in
New York City, New York. He received his Bachelor of
Science degree in Chemistry from the Missouri School of
Mines and] eta11urgy in 1949, From 1941 to 1945 he serv
ed in the United States Army~. From February, 1949 to
August 1950 he was a graduate ~tuaent at the Missouri
School of Mines and etallurgy.
49