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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

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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