United States
Environmental Protection
Agency
Office of Drinking Water
WH550
Criteria and Standards Division
Water
Methods of Removing
Uranium From
Drinking Water
A Literature Survey
Present Municipal Water
Treatment and Potential
Removal Methods
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This document is available to the public through the National Technical
Informntion Service, Springfield, Virginia 22161.
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EPA-570/9-82-003
ORNL/EIS-194
Contract No. W-7405-eng-26
Information Division
METHODS OF REMOVING URANIUM FROM DRINKING WATER:
I. A LITERATURE SURVEY
J. S. Drury
D. Michelson
J. T. Ensminger
Health and Environmental Studies Program
Information Center Complex
II. PRESENT MUNICIPAL WATER TREATMENT AND POTENTIAL REMOVAL METHODS
S. Y. Lee
S. K. White
E. A. Bondietti
Environmental Sciences Division
Work sponsored by the Office of Drinking Water
U.S. Environmental Protection Agency under
Interagency Agreement No. EPA 79-D-X0674
Project Officers
Peter Lassovszky
W. L. Lappenbusch
Date Published: December 1982
OAK RIDGE NATIONAL LABORATORY
Oak Ridge, Tennessee 37830
operated by
UNION CARBIDE CORPORATION
for the
U.S. DEPARTMENT OF ENERGY
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This report was prepared as an account of work sponsored by an
agency of the United States Government. Neither the United States Govern-
ment nor any agency thereof, nor any of their employees, contractors,
subcontractors, or their employees, make any warranty, express or implied,
nor assume any legal liability or responsibility for any third party’s
use or the results of such use of any information, apparatus, product or
process disclosed in this report, nor represent that its use by such
third party would not infringe privately owned rights.
This report has been reviewed by the Office of Drinking Water. U.S.
Environmental Protection Agency, and apprbved for publication. Approval
does not signify that the contents necessarily reflect the views and
policies of the U.S. Environmental Protection Agency, nor does mention
of trade names or coimnercial products constitute endorsement or recom-
mendation for use.
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METHODS OF REMOVING URANIUM FROM DRINKING WATER:
I. A LITERATURE SURVEY
J. S. Drury
D. Michelson
J. T. Ensminger
Health and Environmental Studies Program
Information Center Complex
I—ui
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CONTENTS
Tables I—vu
Abstract 1—ix
1. Introduction 11
2. Removal of Uranium from Ores or Industrial Process
Solutions, Industrial Wastewaters, and Mine Effluents . . . 1—2
3. Removal of Uranium from Seawater I—i
4. Removal of Uranium from Aqueous Analytical Chemistry
Solutions 1—9
5. Removal of Uranium from Natural Fresh Waters 1—12
6. Recommendations 1—19
7. References 123
Appendix 1 1—31
Appendix 2 1—35
I-v
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TABLES
1. Pertinent characteristics of selection systems for removing
uranium from industrial processing solutions, analytical
chemistry solutions, and seawater . . . . . 1—3
2. Uranium in water from U.S. municipal water treatment
plants 1—15
3. Cities from which composite drinking water samples are cur-
rently taken for uranium analyses 1—20
4. Municipal plants producing treated water with relatively
high concentrations of uranium 1—21
5. Commercial and industrial water treatment equipment vendors
and users contacted during this study 1—37
1—vu
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ABSTRACT
Literature was searched for methods of removing uranium from drinking
water. No relevant papers were found, but approximately 1000 publications
were identified in a less specific search for methods of removing uranium
from water. Most of the latter publications dealt with the recovery of
uranium from ores, industrial and analytical chemistry solutions, or sea-
water. The conditions under which these studies were performed were
usually quite different from those normally occurring in municipal water
treatment practice, but some potentially interesting systems of recovery
were identified. A few papers addressed the problem of removing uranium
from natural fresh waters and established the effectiveness of using ad—
sorbents or coprecipitants, such as aluminum hydroxide, ferric hydroxide,
activated carbon, and ion exchangers, under certain conditions. Also,
many U.S. manufacturers and users of water treatment equipment and prod-
ucts were contacted regarding recommended methods of removing uranium
from potable water. Based on the results of these surveys, it is recom-
mended that untreated, partially treated, and finished water samples from
municipal water treatment facilities be analyzed to determine their extent
of removal of uranium by presently used procedures. In addition, labora-
tory studies are suggested to determine what changes, if any, are needed
to maximize the effectiveness of treatments that are already in use in
existing water treatment plants.
I— ix
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1. INTRODUCTION
This survey was prepared by members of the Health and Environmental
Studies Program (HESP), Information Division, Oak Ridge National Labora-
tory, for the Office of Drinking Water (ODW), U.S. Environmental Protection
Agency, to identify technically and economically feasible methods of remov-
ing uranium from drinking water and to indicate problem areas requiring
additional investigation. The survey is part of a larger ODW program to
determine the occurrence of uranium in U.S. surface and groundwaters.
This report is based on extensive contacts with suppliers of water treat-
ment equipment and products and on computer searches of the following data
bases: Chemical Abstracts (1972—80), Nuclear Science Index (1967—76),
National Uranium Resource Evaluation (1945—79), Environmental Aspects of
TransuranicS (1945—78), Chemical Industry Notes (1974—80), Conipendex
(1970—80), Metadex (1966—79), and Current Research (1978—79).
Searches of the above data bases were unproductive when the search
strategy limited the scope of the search to “drinking” water, but approx-
imately 1000 publications were identified when the less restrictive search
descriptor, “water,” was used. In general, most of these publications
concerned the recovery of urani im from water under very special circum-
stances, such as the extraction of uranium from ores or industrial proces-
sing solutions and wastewaters, from seawater, or from solutions used in
analytical chemistry. The greater portion of the cited publications was
not directly applicable to the present study, because boundary conditions
and objectives were drastically different from conditions in potable water
systems. Pertinent characteristics of many of these systems have been
summarized in Table 1, and the most interesting methods are briefly dis-
cussed in Sects. 2, 3, and 4. Techniques deemed more attractive for use
in potable water systems are discussed in Sect. 5, and recommendations
for future research are made in Sect. 6.
I—i
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1—2
2. REMOVAL OF URANIUM FROM ORES OR INDUSTRIAL PROCESS SOLUTIONS,
INDUSTRIAL WASTEWATERS, AND MINE EFFLUENTS
The advent of the Atomic Age in the early 1940s provided the impetus
for the hitherto neglected study of uranium chemistry In general and for
the development of specific methods for recovering uranium from aqueous
solutions. During the subsequent three decades, hundreds of papers were
published describing precipitation, solvent extraction, and ion exchange
techniques applicable to uranium—rich process streams, uranium—depleted
waste streams, and uranium—mine effluents. Many of these papers were
described at one of the three International Conferences on the Peaceful
Uses of Atomic Energy held in Geneva between 1955 and 1964 (Bruce et al.,
1956; Bruce, Fletcher, and Hyman, 1958; Bruce, Fletcher, and Hyman, 1961;
Stevenson, Mason, and Gresky, 1970). More recently, uranium has been
removed from aqueous process streams, wastewaters, and mine effluents by
countercurrent ion exchange (George and Rosenbaum, 1970; Ross and George,
1971), ion exchange membranes (Davis, Wu, and Baker, 1971), organic phos-
phates (Peppard et al., 1973; Tsujino et al., 1970), trioctyiphosphine
oxide (Konstantinova and Mareva, 1977), polyacrylamidosalicycllc acid
(Kennedy et al., 1973), alkyl ammonium compounds (Sato, Kotani, and Good,
1974; Juenger and Schmidt, 1974), ion flotation (Jude and Fratila, 1974;
Zivanov, Miskovic, and Karlovic, 1977), bone char (Blane and Murphy, 1976),
various types of coal (Cameron and Leclair, 1975), and biological sorb—
ents such as fungi, yeasts, and bacteria (Jilek, Fuska, and Nemec, 1978;
Shumate, Strandberg, and Parrott, 1979). Typical examples of the systems
listed above are characterized in Table 1.
It can be seen from Table 1 that most of the listed adsorption,
solvent extraction, and ion exchange systems removed 90% or more of the
initial uranium from the processed solutions. In many Instances, however,
the initial uranium concentrations were so great that even processes having
removal efficiencies greater than 90% still produced effluents containing
as much as 1000 ppb uranium. Concentrations of this degree are unaccept-
ably high for drinking water. If these processes were operated with feed-
stocks containing much lower concentrations of uranium (similar to those
of certain western municipal water treatment plants), some uranium removal
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Table I Pertinent ehn.ctenulwa of aelecled otema Fe. mnnenio uennnnn from induataid ptoeern5 nelulinna. analyneal eheiniaby ndail.ono. .nd minIm
Operatinp tendulern ’
I flTictanl reiyent cit a im jj —
P ineto Na Tcuxarctv aitirtunin ni uranium
me prohiem i toI,,eehuatinn in tilluent
( t ue / I I liii /1
i s ’ fl 0 0 “9 4
I . ‘nm 4
I I iui):’% — tilt)
(natal Initial
reai.ent
con,, ntiutiiun
6 %
ha m mer
hiecnnl.munutii ’n
re.nnne d
l i i
(‘1)
Comn,tnli
Reference
Adtm.rp iuu.n
Li nrre
No
Came,on and LoCh,, 1975
Camcaon and trCIwr. 1175
Cirk e
No
075 n y /mI V aocnuh,c acid added
Van Der cloot dunn and
(harcoal
No
iii aer/ra )
,Iill 94
before enlanciton
Dii. 1975
Shomate Sttandbnrg. and
I’eer,Iounvcen nmeqiienth’
No
I 2 flI
liv
20
I
‘ I S
Panrut 979
5humaln SIrandbor and
. cecc6ernnreetrcrnnunee
No
I 2 nI
i k e
20
( l i i i)
Pmnol. 1979
Shumahe Strandbetp. and
aadn,nrn,esonurnou
No
7 2 p/I
,h
30
II,i mfl
97
>1000 aceanalalino
Panot. 1979
Hnide and Wonenen. t973
Umenflula, flean mImic
Mo
Ii 191k
• 1 )
I act ’ ‘a
Day funpi man contained
Snub patent 1507.001.
Aqierieathae o l g a
Silica pci
No
No
ii (Ilk
a 00 1
90 .100
Ilyb
10-d OmgU/1i 9
1979
PuWal and Sdiwndtau.
1971
Itene 1977
Ilndroualltanlum nude
Tuanium hydannude
No
f in
9
1) 1 10 1
0( 103
31 70
90-.95
Damon. Kemiedy. MdIroy.
and Spance. 964 ire also
Kaatno, l977)
Knyanaka. 1979
Galena
No
9
000)
90
Shretoml. Koimnin, and
Piiyacrylamide gel containing
btantam hydrn’iide
No
3 o
000*
4
66
30
Shtm.pma. 197 7
RtaneandMuaplsy. 1976
Bonn thai
No
61 75
24
Onare. Mutala. Yaanaddha.
Composite Irydroos
blannimliVI muon lIl) nnide
Titanmmlanmnae curnrponnd
No
Pmnbly
40 mg/I
2 vigIl.
6
9
0010
0001
29
Snawulet
and Nahoima. 1979
Maanla. T. 1974
Wootmi. N)r.pnma. Pinion.
I teetunlynin
Mn nemam hydroxide
No
0001
mg Wy
detnioled
unayneoum
hydroxide
end Pallon. 1976
Sch Uss. 1975
I anlb flntnbcrn
Ainenanohil Zepharammn
rcroc hydroxide
Pnmibly
No
01 m M
6 7
0001
0004-00)2
99
92
91
Sodium t hoden ’) taunts
•dded before extraction
SodmmdndecylOatlateadded
Kim and Zontbn, 1971
Willtanto.edGlflun 1979
llydraindtuionnuinnnidn
Na
66
97-95
bnInm nntn.cbon
iudn and floats 1974
Amino rofleat fully arudn
Thnnuinhpdrnoide
Dimethyldoleacyl ammonmm
chloride
No
No
04mM
42 mg/I
3— 9
57
7
0004-0012
50
90
lip),
Sodium dodeconate added
Leunj. Kim ondZntlnn. 1972
Zannoe. MIlome. and
tulane I 77
Dana, Wa . and Bukea. 1971
Inn enc lranne membrane
AMrinoC-103
No
3940
30 000
96
turbo and tier, 1977
km e chonp’ rcoo
I anctmonal pinap II?-
No
1
0010
99
hednonyplaeoylarol 2 naptitol
Arnonicacid
No
5
03
—100
O.OIMEDTAoddedbeIote
entertain
l’nta.odMoyeti. 1976
Jmephuaon. 9976
RA4 I m
No
Alkaline
1000
11)00
975 99
RomandGentpe. 1971
Sirnnnlmneonmnenchahger
S(rimnilnnenniirnnnclianree
No
No
Aikahee
9 I !
00 25—100
l O l l
—0
—IO n
Aococbmceetdandpoenadion
thtocyonate added to complex
uranium
Xn,hlaehandGodl. 974
I - I
1
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Table I (eeahnued)
Enoacemit. teapot. or rena
Ope,aw con00imie
Coamansta
tef a senes
Initial
ivegcnt
CWiCeIIUaIIa%
InutmI
pH
loud
atrium
cmnaetratian
ImaJL)
Conwiunlain
of uranium
in efliaset
taiglL)
titanium
jmmoad
1%)
Decontaraunatiao
(actor
Preicom
Name
Toxicity
problem ’
tiqaadaiaceeebange Aetbeilaim IA•I No 4% in mylene 25-40 511 991 0001K maleoac acid added Dm1” and Kbeç*II.’ ‘7$
(Hdodecyl(tibnethyl)aimee j
Pmtsal pzeaputaooe Smdaaa hydromade No 33 atpjl. $ 0006 6 8 Hadga. Hoffman. Feimnan
and Fcbic. 1914
Hydrcgenndfida Paeaibly IOOmg/L 6-63 50 100 100 999 Koclno,.toeob
Dabinijimb . and
Medve.s. 1971
Renre oonomim Cellelacacetate mentbtmie No 100-0600 98-994 2SOpmjg. ummiyl aBate .olataoua Sau l and Aabbroek. 1976
No 30 <500 250 pug ayntbcoc nn main Saamai and Aâbraeb. 1976
Sohmntexuacaon Ahqaal-336 Na OlMindreUryl O2MNQ 0003 —U — 100 3sMhlhinmcblatldeitdded Barbaaoaadkagm$,I97$
lmneeim
Polyiawa’iomW manor Pimeably 01 mmal m 0 0003 36 TahitI, ICebabe, ad
cyCle Imaabeiaim 5 0 ml cbimofonn Niabipa, 1979
Capiylehydmonamic Ponmbly 05% wfn us 8 10 —100 Venom and Kimena,
acid I .hmaaol 1930
N,N-Di.ai.acrplaee larmde Pamibly I K in cldoroloim 9 K HO 718 lorgaaicjnqanoaa Pebimid l nod Fnlg, 1979
dotnbatmn into,)
Th-e-biilyi Phoitiliale Poambly 99 Tnojao, kacla, Hadano
eadblamam, 1970
Pnobly 0003 —100 Smitten, 1953
Thoctylpheç$ean Pamably 4 7-53 93 tanataudnora and
oxide Mama, 1977
Feb’14. mid Soauikuiido Pambip I 3-90 160-1000 — 100 Kammdy, Saber, NIcal,
mlocydac aced) mid Itaatm, 1973
TncapeylmelliyFamanonaun Pomaibly 002Km kazoos 015 K 700 I I loipnic/aqaeeua 05 K potamiam thmcyaaate Solo, Karma, and Good,
da lmith HO dualoibaima iaiio) added Urmaum amumaceed 1974
aa ananyl Onocymmtn caunphix
Oil l’eihyfleexyftphawbomac Poambly 05 K in Amaco ISO 5-6 K 140-190 7-9 5 93 Haiti, Craiim, and Brown
acid p1w inoclyl- H 3 P0 , 1972
piiaqiiunrie onuda
tjlira liliaauian Hollow nba membomma No Good inc ,a mnded me colloidal, Robs mid Koaaam, 1977
bat not (or drd aranmam
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1—5
would still probably occur, but decontamination efficiencies would
change and laboratory studies based on the new conditions would be
needed to determine the technical feasibility of the processes.
It should also be noted here that a process can remove uranium
efficiently and still not be a practical or economical municipal water
treatment procedure. In addition to the efficiency of the removal of
uranium, the major considerations are capital and operating costs (Sorg
and Logsdon, 1980). If a new treatment process can be added to an
existing plant with minimum modifications or changes in operation, only
small increases in capital and operating costs may occur. If a new
treatment process cannot fit into an existing system without major
changes, substantial capital costs will be incurred in constructing a
new facility, and increased costs may be incurred in its operation.
Based on these considerations, adsorption processes are the most inter-
esting of all the uranium—removing techniques listed in Sect. 2 of this
report. Virtually all municipal water treatment facilities that process
surface waters already have in use the contactors and filters needed in
the removal of uranium by the adsorption technique, and necessary modi-
fications may be effected with minimum incremental cost. However,
municipal water treatment plants that process groundwaters frequently do
not have such facilities; nevertheless, the equipment is commercially
available, and the general technology is well established. Coal, coke,
lignite, bone char, and hydrated titanium oxide or hydroxide are poten-
tially interesting as uranium—removing adsorbents, but data concerning
their uranium—removing efficiencies under typical municipal water plant
operating conditions are needed before a critical selection can be made.
Ion exchange processes constitute the second most interesting group
of uranium-removing techniques listed in Sect. 2. Although it is not
evident from the data in Table 1, there is no technological difficulty
in reducing uranium in processed water to any desired concentration by
treatment with appropriate ion exchange resins (Higgins, 1980). In acid
waters where uranium exists as the uranyl ion, cation exchangers must be
used; for natural waters containing carbonate, where uranium occurs as
the [ U0 2 (C0 3 ) ] complex, anion exchangers are needed. However, impuri-
ties in the processed water can interfere with the intended operation of
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1—6
ion exchange systems, loading the resin with an undesired adsorbent, de-
creasing resin capacity for uranium, and increasing the cost of regener-
ating the resin. In addition, the use of Ion exchangers in the treatment
of municipal waters usually involves relatively high capital costs. For
these reasons ion exchange is used only rarely in municipal water treat-
ment systems that process surface water (Sorg and Logsdon, 1980). In
groundwater treatment plants, cation and anion exchangers are sometimes
used to remove hardness (calcium and magnesium ions) and nitrates. Whether
or not ion exchange techniques should be incorporated into municipal water
treatment plant procedures f or the purpose of removing uranium is thus
primarily an economic question that depends to a considerable extent on
the composition of the water to be treated.
The solvent extraction techniques listed in Sect. 2 and Table 1
generally have high uranium extraction efficiencies and can be used to
deplete uranium in certain effluents to very low levels. However, as a
class, solvent extraction tnethods ire not well suited to the removal of
uranium from drinking water. In general, these two—phase, liquid—liquid
extraction systems require specialized contacting equipment that is not
normally found in conventional municipal water treatment systems. Fur-
thermore, because of the large throughput of most municipal water treat-
ment plants, the contacting equipment must be large and thus relatively
expensive. Although aqueous-organic extraction systems are usually de-
signed for minimum miscibility of the two phases, the aqueous phase
usually contains measurable amounts of the organic solvent or the reagent,
or both. Typically, both reagent and organic solvent are undesirable
drinking water components. Removal of these trace impurities may be
troublesome and expensive. Consequently, solvent extraction techniques
are seldom used in municipal water treatment procedures. Their usage in
the removal of uranium from drinking water should be considered only in
the absence of other more appropriate techniques.
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1—7
3. REMOVAL OF URANIUM FROM SEAWATER
Seawater contains about 3.3 ig/L uranium or an approximate total
of 4.5 billion metric tons (Rodman, Gordon, and Chen, 1979). During the
last 10 years, an anticipated shortage of uranium reserves has stimulated
extensive research into methods of recovering uranium from seawater. This
research is being done primarily in Japan, West Germany, England, and
Russia. Proposed methods include phosphate precipitation (Yamabe and
Takai, 1970), electrochernical deposition (Shigetoini, Kojima, and Shinagawa,
1976; Wooten et al., 1976), anion exchangers (Ryabinin, Lazareva, and
Doroshenko, 1973), and various colloidal flotation schemes (Barannik et
al., 1976; Kim and Zeitlin, 1971; Zhorov et al., 1976). Adsorption of
uranium on various substrates has also been frequently suggested, for
example, on: hydroxylapatite (Takai, Takase, and Yamabe, 1971), surfac-
tants (Ogata and Kakihana, 1969), galena (Koyanaka, 1970; Mukai and
Koyanaka, 1974), peat or brown coal (Astheimer, Schenk, and Schwochau,
1978; Wilhelms, 1972; }leitkamp and Wagener, 1977), macrocyclic hexaketone
(Tabushi, Kobuke, and Nishiya, 1979), benzimidazole polymer (Taniguchi,
Nakayama, and Taril, 1978), hippuric acid formaldehyde copolymer (Taniguchi,
Nakayama, and Tani, 1978), hydrous metal oxides or hydroxides (Ozawa et al.,
1979; Ninomiya, Sugasaka, and Fujii, 1971; Shigetomi, Kojima, and Shinagawa,
1976; Okamoto Yamaguchi, and Takahashi, 1979), silica gel (Ito, Yatnazaki,
and Kantake, 1977; Putral and Schwochau, 1978), chelating resins and fibers
(Egawa and Harada, 1979), quinaldic acid resin (Sakarnoto and Tani, 1977),
and various microorganisms such as algae, fungi, and bacteria (Heide et al.,
1973; Horikoshi, Nakajima, and Sakaguchi, 1979; Jilek et al., 1974). Ad-
ditional papers on methods of removing uranium from seawater are contained
in reviews by Keen (1977), Ogata (1976), Miyazaki (1977), and Novikov and
Lipova (1976) and in bibliographies by Khan and Saleein (1973) and Chen
et al. (1979). Operating conditions and uranium extraction efficiencies
for the most relevant systems listed above are given in Table 1.
Of the various methods proposed for recovering uranium from seawater
there is substantial agreement in the field that retention of dissolved
uranium on a fixed adsorbent is likely to be the most efficient and cost
effective technique. 1-lydrated metal oxides or hydroxides and certain
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1—8
functional groups in the cellular walls of microorganisms appear to offer
potentially attractive adsorptive sites. Among metal hydroxide adsorbents,
hydrous titanium (IV) oxide or titanium hydroxide seems more attractive
than hydroxides of other elements (Liewelyn, 1976; Technology Newsletter,
1980; Kanno, 1977; DavIes et al., 1964). The technological attractiveness
of titanium hydroxide stems from its availability, low unit cost, relative
insolubility in seawater, and high uptake capacity for uranium. Titanium
hydroxide prepared by neutralizing titanium sulfate with an alkali adsorbs
760 iig uranium per gram of titanium, but other methods of preparation, such
as the thermal decomposition of titanium sulfate, can more than double this
uranium uptake (Kanno, 1977). Based on these favorable characteristics,
the Metal Mining Agency of Japan decided to authorize the construction of
a $6.2 million pilot plant based on the use of a titanium hydroxide adsorb-
ent. Construction of the pilot plant, which will be designed to produce
10 kg of uranium per year, will begin in the latter part of 1980 (Technol-
ogy Newsletter, 1980).
Despite the relative advantages of titanium hydroxide over other
metal hydroxide adsorbents in the recovery of uranium from seawater, its
application for this purpose now is not economically competitive with exist-
ing commercial methods of producing uranium (Harrington et al., 1974), nor
is its use expected to contribute significant amounts of uranium to the
world’s stockpile in the near future (Liewelyn, 1976). Koske (1979)
and Rodman, Gordon, and Chen (1979) discuss various design and siting
problems associated with production facilities that use titanium hydrox-
ide adsorbents for recovering uranium from seawater.
Of the various adsorbents listed in this section, peat, coal, hydrous
metal oxides, and microorganisms, such as Paecilomyces marquandii, Sac-
charomyces cerevisiae, and Pseudornonas aeruginosa, appear to have the
greatest potential for removing uranium economically from municipal water
supplies, but more data are needed concerning extraction efficiencies,
loading characteristics, cycling life, and effects of impurities under
typical municipal water treatment plant conditions before any firm recom-
mendation can be made.
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1—9
4. REMOVAL OF URANIUM FROM AQUEOUS ANALYTICAL CHEMISTRY SOLUTtONS
Analytical chemists remove uranium from aqueous solutions by a vari-
ety of techniques, the most important of which are precipitation, solvent
extraction, and ion exchange. In general these methods are similar to
those discussed in Sect. 2 of this report, except that greater emphasis
is attached to reagent selectivity, convenience, or speed and less to
reagent costs and other factors affecting the economics of the process.
Much of the early work in this field is described by Steele and
Taverner (1959), Kraus and Nelson (1956), Peppard (1961), Fans and
Buchanan (1966), Korkisch (1966), O’Laughlin (1966), Morrow (1970), and
Freiser (1970). More recently uranium has been removed from aqueous ana-
lytical chemistry solutions by partial precipitation with alkali (Hodge
et a1., 1974); extraction with trioctyiphosphine oxide (Deutscher and
Mann, 1977), Arsenazo III (Keil, 1979), tributyl phoshate (Gorbushina et
al., 1972), and aimnonium pyrrolidinecarbodithioate (Pradzynski, Henry,
and Draper, 1976); adsorption on chelating resins (Hathaway and James,
1975), anion exchangers (Brits and Smit, 1977; Ryabinin, Lazareva, and
Doroshenko, 1973), silica gel (Putral and Schwochau, 1978), and activated
charcoal (Kuleff and Kostadinov, 1978); electrochemical procedures (Hodge,
1975; Shigetorni, Kojima, and Shinagawa, 1976); various flotation processes
(Barannik et al., 1976; Kim and Zeitlin, 1971; Zhorov et al., 1976); and
reverse osmosis through cellulose acetate membranes (Sastri and Ashbrook,
1976). Operating conditions and extraction efficiencies for typical ana—
lytical chemistry methods are shown in Table 1.
The preceding list of techniques used by analytical chemists to
remove uranium from aqueous solutions includes some potentially interest-
ing methods not previously considered. Partial precipitation with sodium
hydroxide (Hodge et al., 1974) is less effective than other removal methods
( 68% uranium removal from spiked seawater) but is a technique that is
compatible with equipment existing in most present—day municipal water
treatment plants that use surface water sources. The method should be
examined f or effectiveness under municipal water treatment plant conditions.
Reverse osmosis, sometimes called hyperfiltration, can also remove
uranium from water. In the cited literature, initial concentrations of
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1—10
100 to 8000 mg/L uranyl sulfate were reduced 98 to 99.4% using cellulose
acetate membranes and 1786 kPa (250 psi) pressure. Similarly, synthetic
mine effluents containing 30 mg/L uranium were reduced to <500 iigfL. Ura-
nium concentrations of the order of 500 pg/L are unacceptably high for
drinking water, but presumably this concentration could be further reduced
if this were an operational objective. For example, normal operation of
one full—scale reverse osmosis plant reduced the concentration of radium—
226 from 3.4 to 0.26 pCi/L (Sorg and Logsdon, 1980). This 92% removal of
radium—226 was accomplished while producing 1,000,000 gallons per day of
finished water. Obviously, more data are needed to define the efficiency
of reverse osmosis systems with respect to the removal of uranium in the
ppb range, but there appears to be little doubt that the method is tech-
nically capable of lowering the concentration of uranium in treated water.
Whether or not reverse osmosis is economically attractive for this purpose
must be determined on a case by case basis. In general, reverse osmosis
is not economically competitive with conventional municipal water treat-
ment procedures for processing nonsaline feedstocks (Kremen, 1979) but
may have lower combined capital and operating costs than distillation,
freezing, or electrodialysis methods in producing potable water from
brackish water or seawater (Buros, 1979; Larson and Leitner, 1979).
Reverse osmosis units are available in a variety of sizes. Some are
suitable for individual household use (Mgren, 1980).
Another membrane process, electrodialysis, should also receive con-
sideration as a method of removing uranium from drinking water. No lit-
erature reports were found describing the efficiency of this method in
removing uranium from drinking water, but general considerations lead to
the expectation of a decontamination potential similar to that of the
reverse osmosis technique. Electrodialysis systems are commercially
available in a variety of sizes; some electrodialysis units are suitable
for individual household use (Spiegler, 1977). For several years some
highway rest stops and service areas, such as those located at Mohawk
and Sentinel, Arizona, and Junius Ponds, New York, have successfully
used small, unattended, automatic, electrodialysis units to supply
potable water from feedstocks containing high concentrations of calcium
sulfate (Goldstein, 1979; Katz, 1979).
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I—il
Although electrodialysis and reverse osmosis are increasingly used
for desalting seawater and brackish groundwaters, neither process can be
used on turbid surface waters without extensive pretreatment; consequently,
the economic feasibility of using these methods depends on feed water com-
position and must be determined on a case by case basis.
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1—12
5. R. 1OVAL OF URANIUM FROM NATURAL FRESH WATERS
Publications dealing with the removal of uranium from natural fresh
waters have greater relevance to the task at hand than do publications
mentioned in the previous sections of this report, particularly if the
research focused on problems associated with decontamination of water
rather than on the recovery of uranium values. Although few reports have
been published from this perspective, several papers were found that con-
tain pertinent quantitative data.
In the first of these papers, Welssbuch, Cotrau, and Velicescu (1969)
discuss the removal of uranium, cesium—137, strontiuin—90, and yttrium—90
from laboratory prepared aqueous solutions by the addition of varying
amounts of aluminum sulfate and/or activated carbon, followed by mixing,
settling, and filtration through fine porosity filter paper. The change
in concentration of radioactive elements following this treatment was
determined by measuring the initial and final radioactivity of each solu-
tion. Decontamination factors were computed on the basis of reduced
radioactivity. The authors reported removal of 49% of the initial ura-
nium concentration (7.8 tug uranyl nitrate in 100 tuL water) when aluminum
sulfate was added at a rate of 10 m /L. The simultaneous addition of 2
to 3 tng/L of activated charcoal increased the removal of uranium to about
60%. Essentially complete removal of uranium was achieved by the use of
aluminum sulfate alone when the concentration of this flocculant was in-
creased to 40 ing/L (pH not stated). Under laboratory conditions a set-
tling time of one hour was sufficient to achieve effective decontamina-
tion. It should be noted that the initial concentration of uranium in
these test solutions was 1,000 to 30,000 times greater than that normally
encountered in U.S. potable waters, and that the amount of added coagu-
lant varied from about one—half to twice the amounts considered typical
for many U.S. municipal water treatment plants (Durfor and Becker, 1964).
The conditions under which Weissbuch, Cotrau, and Velicescu performed
their experiments are thus not identical to conditions likely to be
encountered in typical U.S. water supplies; nevertheless, it appears
probable that the normal procedure of many U.S. municipal water treatment
plants (i.e., flocculation with alum followed by filtration) removes, or
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1—13
can be readily modified to remove, some or most of the uranium present
in the raw feed water. This observation is at odds with an opinion stated
in an early Water Quality Criteria Report (1968) that coagulation, sedi—
inentation, and rapid sand filtration have little effect on the concentra-
tion of uranyl ions in surface water; consequently, verification of this
conclusion using feed water with more typical uranium concentrations and
closer duplication of standard municipal water treatment plant practices
are desirable.
Laskorin, Metalnikov, and Terentiev (1960) and Laskorin, Metalnikov,
and Sn oliria (1977) also studied the removal of uranium from natural fresh
water. Although their aim was the recovery of uranium rather than the
purification of water, in some respects their experimental conditions
were more pertinent to our task than were those reported by Welssbuch,
Contrau, and Velicescu. For example, Laskorin et al. conducted most of
their experiments with natural lake waters (pH 8.4 to 8.7) containing
about 60 pg/L uranium, a not unusual concentration in some heavily
uranium—contaminated U.S. groundwaters.
Laskorin et al. examined many sorbents and coprecipitants. Calcium
phosphate gel, bone crumbs, Sokolovo bauxite, aluminum hydroxide gel, and
ferric hydroxide gel were found to adsorb uranium; however, the adsorption
capacity of these materials varied greatly. When lake water contained
added uranium (total uranium, 200 ig/L), ferric hydroxide gel had the
greatest capacity for adsorbed uranium (60.1 mg/mL), but this loading
decreased rapidly to only 0.07 mg/mL for lake water containing no added
uranium (total uranium, 60 pg/L). For lake water with no added uranium,
calcium phosphate gel and aluminum hydroxide gel had the best uranium
adsorption capacities, 0.17 and 0.11 tng/mL, respectively. Among synthetic
ion exchange resins, cation exchangers were ineffective, because in lake
water of p 1-I 8.4 to 8.7, uranium existed mainly as the tricarbonato uranyl
anion. However, sorption capacities of 0.18—0.19, 0.24—0.27, and 0.24—
0.26 mg/niL uranium were measured for the anion exchange resins, ED—lOP,
EDE—lO, and A}I—2F, respectively. Data defining minimum concentrations of
uranium in treated water were not provided as the authors were interested
in uranium recovery, not water purification.
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1—14
Among the various methods examined for recovering uranium from lake
water, Laskarin and coworkers preferred coprecipitation with aluminum
hydroxide, followed by settling, filtration, dehydration, and recovery
of the uranium by dissolution in acid. They achieved yields of 80 to
90% using 8 to 24 .ig/L aluminum sulfate. Under the conditions of their
experiments, adsorption of uranium on the flocculant occurred within a
few seconds.
As in the case of Weissbuch and coworkers, the research results
reported by Laskorin et al. are not directly transferable to the puri-
fication of U.S. drinking waters. Nevertheless, their chemical treatments
and process equipment are so similar to those used in typical U.S. water
treatment plants (see Appendix 1) that some removal of uranium must be
presumed to occur, or can be readily made to occur, in present—day U.S.
municipal water treatment plants. Obviously, verification of this pre-
sumption should be possible by comparing the uranium content of raw and
treated waters from various U.S. water purification plants, taking Into
consideration the water sources and the treatment methods employed (Table
2). Unfortunately, no firm conclusions are now possible, since operators
of municipal water treatment plants customarily perform uranium analyses
only on samples of treated, not raw, water. From Table 2 it can be seen
that concentrations of uranium in treated water are generally quite low;
however, without knowledge of the initial uranium concentration in the
untreated water, decontamination factors cannot be determined. Analyses
of the uranium in simultaneously drawn samples of raw, partially treated,
and finished waters from typical U.S. water treatment plants are needed
for this comparison.
-------
2. Uranium in water from U.S. municipal water treatment planta
Table
State and city
Water
aource
Water
treatment
Uranium
concentration
(ppb)
Treated
water
Coagulation,
alum, iron,
salt, or
sludge
Activated
carbon or
silica
Sand or
dintomaceous
aa rt h
filtration
Softening
Chlorination
Fluoridation
Surface 0
crounP
Alabama
Birmingham x x x x <0.1—0.2
Mobile x x x X X 0.1
Montgomery x x x <0.1
Arizona
Phoenix x x x x x x 1.4—5.5
Tucson 1 1 x 1.6—7.3
California
Freano x 0.5
Long Beach x x x x x x <0.1—8.6
Los Angeles x x x x x 4.8
Oakland X x <0.1
Sacramento x x x x x <0.1
SanDiego x x x x x 6.9
San Francisco x x 0.2—0.3
San Jose x x x 0.5
Colorado
Denver x x x x 0.2—2.8
Connecticut
Bridgeport x x <0.1
Hartford x x x x <0.1
New Haven x x x 0.2
District of Columbia
Washington x x x a x <0.1
Florida
Jacksonville i t a <0.1
Miami it a a a a 0.2
St. Petersburg it a a a 0.1
Tampa it it a a a a 0.4
Georgia
Atlanta it a a a <0.1
Savannah a a a a a <0.1
Hawaii
Honolulu it a <0.1
Illinois
Chicago a x a a a x 0.1—0.2
Rockford a 0.6
a
-------
Table 2 (continued)
State
and
city
Water
source
Water
treatment
Uranium
concentration
(ppb)
Treated
water
Coagulation,
alum, iron,
salt, or
sludge
Activated
carbon or
silica
Sand or
diatomaceous
earth
filtration
Softening
Chlorination
Fluoridation
Surf acea
Groundb
Indiana
Evansville X x x x 1 <0.1
Fort Wayne x x x x x X X <0.1
Gary X x x x x x 0.3
Indianapolis x X X X X X 1.3
South Bend X x 0.2—0.3
Iowa
Des Moines x x x x x x x 1.9
Kansas
Kansas City x x x x x 2.4
Topeka X x x x X x <0.1
Wichita x x x x 0.5
Kentucky
Louisville x x x x x x 0.4
Louisiana
Baton Rouge x x <0.1
New Orleans x x x x x <0.1
Shreveport x z x <0.1
Maryland
Baltimore x x x x <0.1
Massachusetts
Boston X X 0.1
Springfield x x x <0.1
Worcester x x <0.1
Michigan
Detroit x x x x x 0.].
Flint x x x x x <0.1
Grand Rapids x x x x x 0.2
Minnesota
Minneapolis x x x x x x x 0.1
St. Paul x x x x x x 0.1
Mississippi
Jackson x x x x x 0.1
Missouri
Kansas City x x x x x X 0.2
St. Louis x x x x x x 0.5
-------
Table 2 (continued)
State
and
city
Water
source
Water
treatment
Uranium
concentration
(ppb)
Treated
water
Coagulation,
alum, Iron,
salt, or
sludge
Activated
carbon or
ciii ca
Sand or
diatoaaceoua
earth
filtration
Softening
Chlorination
Fluoridation
Siarface
crouncS’
Nebraska
Lincoln x x x 5.2
Omaha x x * x x x 2.6
New Jeraey
Jersey City x x 0.1
Newark x x 0.1—0.2
Paterson * x x x x <0.1—0.2
New Mexico
Albuquerque x x 2.9-9.8
New York
Albany * x * * <0.1
Buffalo x * * x 0.2
New York City x * * 0.2—1.0
Rochester * x x * x <0.1
Syracuse x x 0 ,1
Yonkers * x * * 0.1—0.3
North Carolina
Charlotte * * * * * * 0.1
Creensboro x x * x 0.1
Ohio
Akron x x x x * 0.1
Cincinnati x x x x x <0.1
Cleveland x x x x a 0.3
Columbus a * * A * a <0.1
Dayton a x * * <0.1
Toledo a x * a a a a 0.1
Youngstown * a * * a a 0.1
Oklahoma
OkiahomaCity * * a * * a a 0.3
Tulsa a a * a a 0.2
Oregon
Portland a * 0.2
Pennsylvania
Frie * a a a 0.3
Philadelphia * a a a a * <0.1
Pittsburgh a * a a a a a <0.1
-------
Table 2 (continued)
State and city
Water
source
Water
treatment
Uranium
concentration
(ppb)
Treated
water
Coagulation,
alum, iton,
salt, or
sludge
Mtivated
carbon or
silica
Sand or
diatoinaceous
earth
filtration
Softening
Chlorination
Fluoridation
Surfaci
Ground ’
Rhode Island
Providence
x
x
x
x
x
<0.1
Tennessee
Chattanooga
a
x
x
a
0.5
Memphis
a
x
a
<0.1
Nashville
a
a
x
x
a
<0.1
Texas
Amarillo
a
x
4.9—7.4
Austin
a
a
x
X
X
<0.1
Corpus Christi
a
x
a
a
x
a
0.9
Dallas
a
x
a
x
x
x
0.2
El Paso
a
x
a
a
a
a
x
0.1—5.4
Fort Worth
x
a
x
a
1.4
Houston
a
a
a
x
x
a
0.1—2.2
Lubbock
a
a
3.0—250
San Antonio
a
a
0.3
Utah
Salt Lake City
a
a
a
x
a
0.5—2.8
Virginia
Norfolk
a
a
x
a
a
<0.1
Richmond
a
x
a
a
x
a
0.2
Washington
Seattle
a
a
0.1
Spokane
a
X
3.7
Tacoma
a
a
X
0.2
Wisconsin
Madison
a
a
a
0.5
Milwaukee
a
a
a
a
a
<0.1
Surface water sources include rivers, other streams, reservoirs, and lakes or ponds.
Groundwater sources include ordinary and artesian wells and springs.
Source: Adapted from Durfor and Becker, 1964.
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1—19
6. RECOINENDATIONS
The literature discusses many techniques of recovering uranium from
aqueous solutions, but no papers address the specific problem of removing
this pollutant from drinking water. In the few papers that contain use-
ful data, there is evidence that, under favorable conditions, uranium at
concentrations likely to be encountered in potable water sources may be
removed by several processes that are, or could readily become, part of
most municipal water treatment procedures. Some of these processes are
adsorption on aluminum or iron coagulants, activated carbon or silica,
clays, or the cellular walls of microorganisms that compose the schmutz-
decke (see Appendix 1) that coats the sand particles of most sand filters
used in municipal water treatment plants.
Whether or not uranium is removed in typical municipal water treat-
ment procedures should be easy to ascertain; one need only compare the
concentrations of uranium in simultaneously drawn samples of raw and
treated water. Unfortunately, such analytical data for comparable raw
and treated water samples could not be found for any U.S. municipal water
treatment plant. Apparently because of the expense involved, uranium
analyses when performed at all are usually made only on treated water
samples.
Although almost all such published analyses show very low uranium
concentrations, suggesting that uranium removal does indeed occur, it is
not possible to verify this effect without analytical data on comparable
raw water samples. Priority should therefore be given to securing and
analyzing samples of raw and treated water from typical U.S. water treat-
ment plants. Samples should be collected after each step in the treatment
process so that effective decontamination procedures can be identified.
Collection and analysis of such samples could be implemented at relatively
small incremental cost by the U.S. EPA Office of Radiation Programs, Mont-
gomery, Alabama, which already routinely analyzes uranium in treated water
samples from 20 municipal water plants across the United States (Table 3).
Alternatively, water from the cities shown in Table 4 could be sampled.
The waters from these cities contain the highest recorded concentrations
of uranium in municipally treated waters that we encountered in our
survey of the literature.
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1—20
Table 3. Cities from which composite drinking water
samples are currently taken for uranium analyses
Total uranium concentration
Location July-December 1977 composite
(pCi / L)
Baruwell, South Carolina 0.047
Berkeley, California 0.041
Bismarck, North Dakota 0.123
Chicago, Illinois 0.379
Columbia, South Carolina 0.027
Columbus, Ohio 0.065
Denver, Colorado 2.280
Harrisburg, Pennsylvania 0.030
Knoxville, Tennessee 0.062
Las Vegas, Nevada 5.341
Los Angeles, California 2.933
Lynchburg, Virginia 0.065
Miami, Florida 0.181
Montgomery, Alabama 0.044
Niagara Falls, New York 0.263
Oklahoma City, Oklahoma 0.170
Pittsburgh, Pennsylvania 0.141
Portland, Oregon 0.028
Santa Fe, New Mexico 0.303
Trenton, New Jersey 0.037
Source: U.S. Environmental Protection Agency, 1979.
If the results of the study suggested above are inconsistent or in-
conclusive, it is recoended that laboratory studies be initiated under
controlled conditions to determine if existing municipal water treatment
procedures adequately remove, or can be readily modified to adequately
remove, uranium from drinking water. The studies should include conven-
tional water plant coagulants and adsorbents, such as alum, iron salts,
activated carbon and silica, and schmutzdecke, as well as natural and
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1—21
Table 4. Municipal plants producing treated water
with relatively high concentrations of uranium
Location pg/L pCi/L
(computed)
Albuquerque, New Mexico 9.80 6.66
Amarillo, Texas 7.40 5.03
Des Moines, Iowa 1.90 1.29
Denver, Colorado 2.80 1.90
El Paso, Texas 5.40 3.67
Fort Worth, Texas 1.40 0.95
Houston, Texas 3.60 2.45
Kansas City, Missouri 2.40 1.63
Lincoln, Nebraska 5.20 3.54
Long Beach, California 8.60 5.85
Los Angeles, California 4.80 3.26
Lubbock, Texas 13.00—250.00 8.84—170.00
Midland, Michigan 20.00 13.60
Omaha, Nebraska 2.60 1.77
Phoenix, Arizona 5.50 3.74
Salt Lake City, Utah 2.80 1.90
San Diego, California 6.90 4.69
Sault St. Marie, Michigan 37.00 25.16
Spokane, Washington 3.70 2.52
Tucson, Arizona 6.20 4.22
Source: Durfor and Becker, 1964; Scott and Barker,
1962; U.S. Environmental Protection Agency, 1979.
synthetic ion exchangers. Adsorption should be studied as a function of
uranium concentration in raw water, type and concentration of adsorbent,
adsorption pH, adsorption kinetics, and influence of coon cationic and
anionic impurities. The aim of the research should be the adaptation of
existing water treatment practices with minimal economic impact. Studies
are also needed to determine the operating conditions for optimum uranium
removal by the electrodialysis and reverse osmosis processes. Neither
-------
1—22
of these processes is likely to be the method of choice for the removal
of uranium from drinking water in communities with a heavy investment in
conventional water treatment equipment and access to supplies of good
quality fresh water. One of these treatment processes, though, may well
be chosen for use by communities without such supplies of water or by
communities that must expand service beyond the capacity of such existing
supplies. The electrodialysis and reverse osmosis studies should deter-
mine uranium removal as a function of membrane type, membrane preparation
methods, pH, pressure/voltage, and feed water composition, including tur-
bidity and common cationic and anionic impurities.
Attention should also be directed to the development of appropriate
techniques for removing uranium from drinking water that is supplied by
sources other than municipal water treatment plants. It is likely that
the most serious exposures to uranium—polluted water occur among the small
fraction of the population that cotLsumes unprocessed water from private
wells located in uraniferous strata. Although relatively few drinking
water wells are known to contain high concentrations of uranium, water
from some wells contains uranium in excess of 100 ppb and may contain more
than 400 ppb (Wagoner, 1979). Since the required volume of treated water
from private wells is much smaller (minimum for drinking water, 2 L per
person per day) than that from municipal water plants, normal constraints
imposed by capital investment requirements for municipal water plants do
not apply, and techniques not feasible for municipal water plants can be
considered for private wells. Very likely, convenience, rather than oper-
ating cost or capital investment, would be the dominant consideration.
Under these circumstances, ion exchange chromatography, adsorption on
suitable substrates, such as molecular sieves or hydrous metal oxides,
or a membrane process such as electrodialysis or reverse osmosis would
probably be the preferred purification method. Such a choice would prob-
ably allow use of much present—day water softening methodology and equip-
ment. However, loading and extraction efficiency studies would be needed
to define the best types of adsorbents, ion exchange resins, or membranes,
as well as optimum operating conditions for the removal of uranium from
typical U.S. groundwaters. Information of this type is not now available
from either commercial water treatment equipment vendors or industrial
users of this equipment (Appendix 2).
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1—23
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APPENDIX 1
U.S. Municipal Water Treatment Practices
Public water supplies are subjected to a variety of treatments,
depending on the condition of the raw water and the needs of the consum-
ing community. The following discussion is based on the principal steps
involved in treating a hypothetical “very hard” surface water that is
laden with silt (Durf or and Becker, 1964). Few municipalities will employ
all of the indicated processes, but most will utilize several. In most
instances groundwaters undergo less extensive processing than surface
waters. These processes usually include disinfection, ion exchange, or
occasional lime softening. Treatment of surface waters may include some
or all of the following processes dependent on water quality. These
processes may include screening, coagulation, precipitation, filtration,
ion exchange, lime softening, post—chemical treatment, and disinfection.
Following are brief descriptions of these processes.
Screening
Water is pumped through grates to remove tree limbs and other
floating debris.
Coagulation and precipitation
After screening, coagulation and precipitation chemicals are added
to the water to remove sediments, turbidity, color, and organic matter.
Coagulation and precipitation processes change the properties of
dissolved, colloidal, and suspended materials so that contaminants settle
out of solution by gravity. The stability of colloids and suspended or
dissolved solids is due to electrical charge and solvation effects. An
effective coagulant changes the surface charge properties of particles
so that they tend to agglomerate, or it enmeshes particles in a polymeric
suspension that readily settles under the influence of gravity. Commonly
used coagulants include alum [ Al 2 (S0 4 ) 3 .l41-1 2 0), ferric chloride (FeCl 3 ),
ferrous sulfate (FeSO 14 ), lime [ Ca(OH) 2 ], and a variety of polymers and
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clays. An effective precipitant is a chemical that combines with a
dissolved contaminant to form an insoluble compound that settles out of
solution or can be readily removed by filtration. Typical precipitants
include soda ash (Na 2 CO 3 ), caustic soda (NaOH), and potassium perman—
ganate (I 4n0t ).
Coagulants and precipi.tants are usually added to raw water in solid,
slurry, or dissolved form, with 1 to 2 minutes of rapid mixing, followed
by reduced agitation for about 30 minutes to stimulate floc growth.
After the flocs reach optimum size, the treated water passes to primary
and secondary settling basins where most of the solids separate from the
water after a residence time of several hours.
The kind and amounts of coagulants and precipitants added to raw
water vary depending on the nature of the influent water. Raw water
containing less than 50 mg/L total dissolved or suspended solids is
difficult to treat effectively because flocs do not readily form in such
a dispersed system. On the other hand, treatment of water containing
dissolved or suspended solids in excess of 2000 mg/L is also troublesome
because settling of the agglomerated solids is hindered. The pH of the
water is important for effective coagulation. When alum is used, the
most effective coagulation occurs in the pH range of 5 to 7; with lime,
a pH of 7 to 10 is preferred. Following coagulation, some sort of
neutralization is usually performed to return the pH of the treated water
to a value near 7.
Coagulation and precipitation produce considerable quantities of
sludge solids. Typically, alum produces 0.45 kg for each 100 to 200 g
of added aluminum. Lime produces about 450 to 600 kg of sludge per 1000
m 3 of soft water, while the use of anhydrous ferric chloride under
similar circumstances results in about 50 kg of solids. Such large
quantities of sludge present substantial disposal problems. Scaling can
also be a problem when large quantities of solids are present.
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Filtration
Clarified water from the settling basins is typically filtered
through several feet of fine sand or through a layer of anthracite coal
(anthrafilt) to remove the remaining suspended matter and most bacteria.
After several days of use, a freshly cleaned bed of filter sand
develops a slimy deposit of sediment and bacteria, called schmutzdecke,
which may remove bacteria from water (Durf or and Becker, 1964) and may
be effective in reducing the uranium content of treated water (Shumate,
Standberg, and Parrott, 1979). Filters charged with crushed anthrafilt
also have a potential for removing uranium (Cameron and Leclair, 1975).
If the treated water was previously softened by the addition of
lime or lime soda, it is now saturated with calcium carbonate. To
prevent subsequent precipitation of this material in water pipes and
other equipment, it is desirable to convert the calcium carbonate to a
more soluble salt, such as calcium bicarbonate. This is usually accom-
plished by adding sulfuric acid or by injecting carbon dioxide gas. In
some water treatment plants, this stabilization procedure is performed
before the filtration step to prevent clogging of filter beds, but other
water treatment plants stabilize the water after filtration.
Softening
The presence of calcium and magnesium salts is chiefly responsible
for hardness in water. To soften water, the concentration of calcium
and magnesium ions must be reduced or removed. In most municipal water
treatment plants, softening is accomplished by adding lime or lime and
soda ash to the water. The addition of lime or lime and soda ash converts
dissolved calcium and magnesium salts into a sludge of insoluble calcium
carbonate and magnesium hydroxide. This sludge gradually settles and
carries down suspended sediment, bacteria, and finely divided organic
matter, thereby increasing the efficacy of the coagulating chemical.
In a few municipal water treatment plants, calcium and magnesium in water
are replaced with sodium by means of a cation exchange process utilizing
zeolites or synthetic ion exchange resins. This water softening process
is very effective, but the resulting water may be aggressive, requiring
the addition of inhibitors to reduce the corrosivity.
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Taste and odor removal
In a sense each water treatment step contributes to the removal of
tastes and odor from the treated water. However, some water treatment
plants also aerate water to reduce tastes and odors. Some plants may
add activated carbon with coagulation and softening chemicals to adsorb
undesirable tastes, odors, and colors from treated water. Eventually,
this carbon becomes part of the coagulated floc and settles as sludge.
The use of activated carbon in municipal water treatment plants is of
special interest to this study since some research (Weissbuch, Cotrau,
and Velicescu, 1969) indicates it to be effective in removing uranium
from water under certain conditions.
Disinfection
Before distribution and use, the stored finished water is usually
freed from disease germs and other harmful microorganisms by treatment
with chloride, hypochlorite, or chloramine compounds. In some localities,
ozone rather than chlorine is used as a disinfectant.
Prior to the 1970s, sometimes chlorine was added to untreated
influent water immediately after the screening step as a means of con-
trolling the growth of plants and microscopic organisms that could impart
undesirable tastes and odors to the water. This procedure, called pre—
chlorination, generally is not now recommended for most water because of
the high probability of forming toxic and carcinogenic haloforms from
naturally occurring trace organic compounds in the feed water.
Postchemical treatment
If the treated water is to be fluoridated, sodium fluoride, sodium
silicofluoride, or fluorosilic acid is generally added to the water.
This addition is usually performed after the filtration step because
fluoride can be removed by lime—softening and alum—coagulation treatments.
Lime may also be added after filtration if the pH or hardness of the water
is less than desired. After processing, the treated water passes to
storage reservoirs to await distribution to customers.
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1—35
APPENDIX 2
Equipment Vendors and Users Contacted
The vendors or users of commercial water treatment equipment and
supplies listed in Table 5 were canvassed during the preparation of this
report concerning the possible use of their product or system to reduce
uranium in drinking water to the ppb level. Most of those contacted had
no data for uranium removal. None of those contacted had relevant exper-
imental data for the performance of their system product at ppb concen-
trations. In general, economic considerations dictate the recovery of
uranium from wastewaters down to the ppm level, but neither economic nor
legal incentives presently exist for industry to reduce uranium In efflu-
ents to concentrations lower than the ppm level. For example, operators
of uranium mines in Colorado, Wyoming, and the Grants Mineral Belt, New
Mexico (Kermac Nuclear Fuels Corporation, United Nuclear — }Iomestake
Partners, the Anaconda Company, and Union Carbide Corporation), have
National Pollutant Discharge Elimination System permits allowing the dis-
charge of effluents containing 2 (average) or 4 (maximum) mg/L uranium
(Beverly, 1980; Dehn, 1980; Rouse, 1980). Consequently, there has thus
far been no industrial demand for, or commercial development of, processes
that reduce uranium in effluents to the ppb level.
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1—36
Table 5. Commercial and industrial water treatment equipment
vendors and users contacted during this study
Equipment vendor or user Address Contact
Anaconda Company, Mineral
Resources Group
Babcock and Wilcox Company,
Nuclear Fuel Division
Bruner Corporation
Chemical Separations
Corporation
Consolidated Edison
Diamond Shamrock
Corporation, Functional
Polymers Division
Dow Chemical, USA
Mach Chemical Company
Freeport Uranium Recovery
Company
lonac Chemical Company
Nuclear Fuel Services
Nuclear Metals,
Incorporated
Oak Ridge National
Laboratory
Osmonics, Incorporated
Rainsoft Water
conditioning Company
Rohm and Haas
Technic Central Systems,
Incorporated
The Lindsay Company
Union Carbide Corporation,
Metals Division
Water Equipment Technol-
ogies, Incorporated
Watco R.O.
Zeolite Chemical Company
Denver, Colorado
Appollo, Pennsylvania
Milwaukee, Wisconsin
Concord, Tennessee
New York, New York
Cleveland, Ohio
Midland, Michigan
Loveland, Colorado
UnclE Sam, Louisiana
Birmingham, New Jersey
Rockville, Maryland
Concord, Massachusetts
Oak Ridge, Tennessee
Hopkins, Minnesota
Elk Grove Village,
Illinois
Philadelphia, Pennsylvania
Seattle, Washington
St. Paul, Minnesota
Grand Junction, Colorado
Palm Beach, Florida
Las Vegas, Nevada
Clayton, New Jersey
R. McClincy
C. Del Signore
T. Christiinan
I. Higgins
R. Van Wyck
J. Griggs
L. LeFevre
S. Whitmore
J. Jolly
F. McGarvey
R. Idaker
A. Gilman
C. Strandberg
S. Hurt
J. Johnson
L. Comb
A. Graham
B. Robbins
L. Mgren
D. Windberg
R. J. Beverly
L. Stenger
S. Casey
R. Fackler
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METHODS OF REMOVING URANIUM FROM DRINKING WATER:
II. PRESENT MUNICIPAL WATER TREATMENT AND POTENTIAL REMOVAL METHODS
S. ‘1. Lee
S. K. White
E. A. Bondietti
Environmental Sciences Division
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CONTENTS
Figures . I l —v
Tables 11—v i A
Abstract I lix
1. Uranium Removal by Present Municipal Water Treatment
Processes
1.1 Sample Collection
1.2 Sample Analysis
1.3 Results and Discussion
2. Examination of Available Removal Methods
2.1 Uranium Removal by Alkaline Earth Water Softeners
Lime [ Ca(OH) 2 ) Treatment
Ca(OH) 2 and MgCO 3 Treatment
pH and Uranium Concentration Effects
Discussion
2.2 Uranium Removal by Coagulants
2.2.1 Fe 2 (SOi,) 3 Treatment .
2.2.2 FeSO 4 Treatment
2.2.3 Al 2 (SO ) 3 Treatment .
2.2.4 Discussion
2.3 Uranium Removal by Absorbents
2.3.1 Titanium Oxide (TiO 2 ) . . *
2.3.2 Activated Charcoal .
2.3.3 Discussion
2.4 Uranium Removal by Ion Exchangers
2.4.1 Batch Test
2.4.2 Column Test
2.4.3 Discussion
2.5 Uranium Removal by Reverse Osmosis
3. Summary of Results . . . . 11—44
4. Recommendations 11—46
5. References 11—47
Appendix . 11—49
11—1
1 1—1
11—2
1 1—2
2.1.1
2.1.2
2.1.3
2.1.4
11—10
11—12
11—12
1 1— 13
11—14
11—16
11—20
11—21
11—21
11—22
11—23
11—25
11—26
11—26
* 11—28
11—29
11—29
11—31
11—40
11—42
Il—Hi
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FIGURES
1. Percent uranium (initial uranium concentration, 48 gIL)
removed from 5 x iO M ?4gCl 2 , 5 x lo M MgCl 2 5 x iO — M
Na}1C0 3 , and 5 x iO M CaC1 2 —5 x M NaHCO 3 solutions
varying pHs 1 1]8
2. Percent uranium (initial uranium concentration, 48 pg/L)
removed from 5 x iO M MgC1 2 and from the mixture of
5 x i0 3 M MgC1 2 —5 x o- 3 M NaHCO 3 solution after addition
of varying amounts of 1 M NaOH solution 11—19
3. Percent uranium removed from pond water (initial uranium
concentraion, 83 ug/L) as a function of doses (vigiL) of
Fe 2 (S0 4 ) 3 and pH 11—22
4. Percent uranium removed from pond water (initial uranium
concentration, 83 ig/L) as a function of doses (vigiL) of
FeSO and pH 11—23
5. Percent uranium removed from pond water (initial uranium
concentration, 83 iig/L) as a function of doses (vigiL) of
Al 2 (S0 4 ) 3 and pH 11—24
6. Percent uranium (initial uranium concentration, 23.8 vigiL)
passed through an anion exchange column versus cumulated
column volumes of inf].uent solution 11—38
7. Percent of uranium eluted from uranium—containing anion
exchange column (initial uranium loading, 39 mg) by 1.0 M
NaCl—0.5 M NaHCO 3 solution versus cumulated column volumes
of elutiori solution 11—39
Il-v
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TABLES
1. Uranium concentration in raw water, intermediate stage, and
product water samples taken from selected municipal water
plants . 11—3
2. Alkalinity versus uranium content of the municipal water. . 11—7
3. Effect of raising the pH followed by filtration (0.22 inn)
on the uranium concentration in water from U.S. municipal
treatment plants 11—9
4. Chemical composition of pond (3513) water 1 1-11
5. Removal of uranium from pond water by Ca(0H) 2 treatment . . . 11-13
6. Removal of uranium from pond water by combined Ca(OH) 2
and MgCO 3 treatment 11—14
7. Effect of pH at a given Ca(OH) 2 and MgCO 3 dose on uranium
removal from pond water 1115
8. Efficiency of uranium removal by Ca(OH) 2 and MgCO 3 treat-
ment at varying uranium concentrations 11—16
9. Percent uranium removal by Fe 2 (S0i 4 )3, FeSOi. 1 ,, and A1 2 (SOi..) 3
coagulants with varying pH 11—21
10. Suspected uranyl species and charge characteristics of iron
and aluminum hydroxide flocs at given pHs of pond
water 11—25
11. Effects of carbonate concentration and pH on the adsorption
of uranium on titanium oxide 11—27
12. Effects of carbonate concentration and pH on the adsorption
of uranium on activated charcoal 11—28
13. Adsorption of uranium by anion exchange resin from waste
pond vater at varying pH 11-30
14. Adsorption of uranium by anion exchange resin from 0.01 M
CaSOj. 4 solution at varying pH 11—31
15. Adsorption of uranium by anion exchange resin from 0.001 M
CaSO 4 solution at varying pH 11—32
16. Adsorption of uranium by anion exchange resin from 0.02 M
NaC1 solution at varying pH 11-33
17. Removal of uranium by Ca— and Na—resin column at varying
solution pH 11—34
18. Percent uranium recovered from cation resin columns by
various eluding solutions 11—35
19. Uranium adsorption by two different anion exchange columns
with varying flow rate of 237 U—spiked pond water 11-36
h—vu
-------
20. Removal of uranium by a O.5—mL anion resin column 11—37
21. Percent of uranium loading on resin column at the selected
uranium concentration in the effluent 11—39
22. Percent 237 U recovered from anion resin columns by various
eluting solutions 11—40
23. Percent of uranium removed by a reverse osmosis module . . . 11—43
II—viii
-------
ABSTRACT
Uranium analyses of raw water, intermediate stage, and treated
water samples from 20 municipal water treatment plants indicated that
the present treatment practices were not effective in removing uranium
from raw waters when the influent concentration was in the range of 0.1
to 16 pg/L uranium. Laboratory batch tests revealed that the water
softening and coagulant chemicals commonly used were able to remove more
than 90% of the dissolved uranium (<100 zg/L) in waters if an optimum
pH and dosage were provided. Absorbents, titanium oxide and activated
charcoal, were also effective in uranium removal under specific condi-
tions. Strong base anion exchange resin was the most efficient uranium
adsorbent, and an anion exchange column is a recommended option f or the
treatment of private well waters containing uranium at higher than
desirable levels.
IT—ix
-------
1. URANIUM REMOVAL BY PRESENT MUNICIPAL WATER TREAThENT PROCESSES
The effectiveness of water treatment processes for the removal of
inorganic contaminants (including radionuclides) has been su narized by
Sorg and Logsdon (1980). However, very little information is available
on the effectiveness of current treatment processes in removing uranium.
As part of an interagency agreement between the Office of Drinking Water
(ODW), the U.S. Environmental Protection Agency (EPA), and the U.S.
Department of Energy (DOE), a study of the efficiency of uranium removal
in water processed at municipal treatment plants was conducted. Water
from 20 municipal water purification plants representing eight states
was studied. The municipal plants sampled were chosen by EPA because
noticeably higher concentrations of uranium in previous measurements of
treated water had been shown there. The study was designed to determine
whether existing water treatment practice can remove uranium to safe
levels for drinking waters and to investigate variations or additions to
water treatment methods for removing the uranium to safe levels.
1.1 SAMPLE COLLECTION
Collapsible polyethylene one—gallon containers with screw caps and
shipping cartons were mailed to the sampling personnel at selected
municipal water treatment plants or to the EPA regional staff for col-
lection of the water samples. Each sampler received one set of four
containers with cartons for each plant and was instructed to collect a
raw water sample, a water sample after the coagulation and/or softening
step, if applicable, and a treated (end—product) water sample. An extra
container was enclosed to be used at the sampler’s discretion for
sampling the water during another stage of the treatment process. A
form letter was enclosed with each set of containers requesting the date
of sampling, the water source, and descriptive remarks or measurements
regarding the samples taken. The samples received were stored at room
temperature, with the exception of the raw water samples which were kept
in the refrigerator (4°C).
11 —1
-------
11—2
1.2 SAMPLE ANALYSIS
A raw water sample, a water sample from intermediate stages, and a
final treated water sample were analyzed for uranium by neutron activa-
tion analysis after concentration of the uranium from the sample solu-
tion onto an anion exchange column. The experimental method was similar
to that used by Brits and Sinit (1977). Optimum conditions for adsorption
of the uranium from auunonium carbonate systems with methanol were pre-
viously determined by Haggag and Stokely (1981). They reported that
greater than 99% of the uranium was adsorbed on the column and that
quantities of uranium from 0.01 to 6.6 mg (in 100-niL solution) were re-
tained by the column from a O.1—M ainmonium carbonate—methanol system.
The efficiency of the method was determined with 237 U, a gamma emitter.
Controls of deionized water were run at intervals to ensure that the
equipment remained free of contamination. Detailed procedures for sample
preparation and uranium analysis are presented in the appendix.
1.3 RESULTS AND DISCUSSION
The results show that very little uranium is removed by present
water treatment methods. Table 1 summarizes the results of the uranium
analyses of the raw water, intermediate stage, and treated water samples
and describes the point in the treatment process at which each sample
was collected. Uranium concentrations are reported in micrograms per
liter ( ig/L), and their range is from the minimum detectable concentra-
tion (0.01 ig/L) to 17 .ig/L. Measurements of pH are also recorded in
Table 1. The lowest pH reported was 7.4, and the highest was 9.5. In
some instances, the uranium level in the treated water was higher than
that in the raw water sample. This probably resulted from blending
waters from two different sources, one with a higher uranium level than
the other. Most of the plants used a combination of conventional
methods — coagulation (alum, lime, iron salts, or polymers), sedimenta-
tion, filtration, and chlorination. Very few of the plants employed
activated carbon for taste and odor control.
An interesting case in which the uranium analyses suggested that some
removal of uranium may have taken place was in the samples from Kansas City,
-------
Table 1. Uranium concentrations in raw water, intermediate stage, and product water samples taken from
selected municipal water plants
Analyses
Location plant Stage Datea Source/Description pH 1 ’ U,
Midland, Mich. Raw 12—4—80 1 Lake Huron, before Cl 2 added 7.7 0.27 ± 0.03, n = 3
2 After Cl 2 added 8.0 0.47
Intermediate After Fe 2 (S0t 4 ) 3 coagulation 8.1(7.9) 0.59
before sand (iltration
Final 7.8 0.35
Houston, Texas Raw 11—20—80 San Jacinto River, taken from 7.4 0.32 ± 0.14, n = 3
holding pond
Intermediate 1 After alum coagulation 7.6 0.37
2 Supernatant 0 iter lime 7.9 0.28
stabilization
Final 7.7(1.0) 0.28
Kansas City, Mo. Raw 12—18—80 Missouri River 8.1(7.9) 5.33 ± 0.23, n 3
Intermediate 1 After addition of lime plus 8.0(8.6) 4.86
Cat—Floc C polymer
2 After addition of Fe 2 (S0i ) 3 , 9.4(9.9) 4.73
Cl 2 , NH 3 , and Ca(OH) 2
Final After sand filtration, Cl 2 9.5(9.9) 4.07 ± 0.30, n = 4
disinfection
Lincoln, Neb. Raw 12—16—80 Well water 8.0(7.4) 7.29 ± 0.47, n 3
Intermediate 1 Filter influent, Cl 2 , NR 3 ,and 8.3 9.11
H 2 SIF 6 added
2 Filter effluent 8.4 8.91
Final Pump discharge 8.0 7.39
Denver, Cob. — Raw 1—14—81 1 Blue River & So. Platte River, 7.6 1.60, 1.60
Marston at prechiorination
2 7.6 1.50
Intermediate After polymer coagulation plus alum 7.8 1.50
Final After filtration 7.8 1.50
-------
Table 1 (continued)
Analyses
Location plant Stage Datea Source/Description U, pg/L 0
Mof fat Raw 2—5—81 Winter Park Ralston Reservoir 7.5(7.5) 15.9 ± 1.58
Intermediate 1 After alum and iron polymer 7.2 3.90
coagulation
2 After lime addition 7.8(8.3) 4.90
Final 7.6(7.3) 4.00
San Diego, Calif. — Raw 12—26—80 1 Local water (75%) 7.8(7.9) 5.43, 5.10
Escondido Vista 2—Calif. aqueduct (25%) 8.1(7.7) 8.15 ± 0.39, n = 3
Intermediate After coagulation with alum 7 .7 6.53
(blend) and cationic polymer
Final (blend) After filtration 7.7(7.6) 6.25, 5.60
Alvarado Raw 12—26—80 San Vicente and El Capitan 7.8(1.7) 1.68, 1.90
reservoirs H
H
Intermediate After coagulation with FeC1 3 , 8.2(8.2) 2.12
lime, CaO
Final After Cl 2 disinfection 8.2(8.2) 2.31, 1.80
Otay Raw 12—26—80 Barrett Reservoir 7.9(8.0) 1.05, 1.00
Intermediate After coagulation with FeCI 3 7.7 1.22
Final After filtration and Cl 2 8.2(8.2) 2.10, 2.30
disinfection
Sweetwater Raw 12—26—80 Sweetwater Lake 8.0(8.1) 4.07, 3.30
Intermediate Filter influent after coagula— 8.0(7.7) 4.25
tion with cationic polymer
Final After filtration and Cl 2 8.1(7.8) 3.30
disinfection
Los Angeles, Raw 12—11—80 Well water plus Cl 2 7.7(8.0) 0.14, 0.29, <0.10
Calif. — Intermediate 1 After lime addition, before 8.0 0.43
Hawthorne sedimentation
2 After sedimentation 7.9 0.30
Final After filtration 8.2 0.10, <0.10
-------
Table 1 (continued)
Analyses
Location plant Stage Datea Source/Description pHi’ U,
Jensen — Raw 2—9—81 State project water 7.8 0.30, 0.27
Metropolitan Intermediate 1 Settled water after alum coagu- 8.0 0.25
Water District latlon
2 Clarified water after filtration 8.2 0.27
Final 8.2 0.28
Long Beach Raw 12—9—80 Well water 8.2 0.32
Intermediate 1 After lime coagulation plus Cl 2 8.1 0.65
2 After filtration 8.2(7.7) 0.51
Final After blending with treated 8.3(7.6) 1.59 ± 0.33, n = 3
surface water
Weymouth — Raw 12—4—80 Colorado River 8.0 6.61, 6.10
Metropolitan Intermediate 1 Flocculator effluent 8.0 7.51
Water District 2 Sedimentation basin outlet 7.9 7.40 ‘-4
Final 8.0 6.06, 6.60
Phoenix, Arjz. — Raw 12—10—80 Verde River 8.4 4.45, 4.30
Val Vista Intermediate 1 After coagulation with alum 8.1 4.76
2 Presedimentation before chemical 8.4 4.50
addition
Final After filtration and chlorination, 8.2 4.05
activated carbon (dual media)
Verde Raw 12—10—80 Verde River 8.3 4.30, 4.10
Intermediate 1 After coagulation with alum 8.2 6.20
2 Filter influent 8.2 4.32
Final After rapid sand filtration and 8.2 4.20
chlorination, activated carbon
Squaw Peak Raw 1—20—81 Salt RIver 8.2 1.70
Intermediate 1 From sedimentation basin after 7.9 1.40
coagulation with alum
2 Effluent from storage 8.2 1.50
Final After sand filtration, Cl 2 , actl— 8.2 1.80
vated carbon
-------
Table 1 (continued)
aDate of sampling.
bpH recorded at time of analysis; pH reported by water sampler in parentheses.
CFor replicate analyses (n) greater than 2, values are mean ± 1 standard deviation.
Location plant
Stage
Datea
Source/Description
Analyses
pHi’
U,
Salt Lake City, Utah
Raw
2—25—81
City Creek
7.9
1.00,
0.80
City Creek
Intermediate
1 Collected at end of flocculation
basin
2 Collected at end of sedimentation
basin
1.00,
0.90,
1.00
0.90
Final
0.90,
1.00
Little Cottonwood
Raw (blend)
Intermediate
Final
3—3—81
Little Cottonwood Creek (22%)
Deer Creek Reservoir (78%)
1 After coagulation with alum
2 After sedimentation
7.9
1.70
1.50,
1.30,
0.90,
1.50
1.60
1.00
Big Cottonwood
Raw
Intermediate
2—24—81
Big Cottonwood Creek
1 Collected at end of flocculation
basin
2 Collected at end of sedimentation
basin
8.3(8.1)
8.0
8.0
0.90
0.80
0.80
Final
8.1
0.80
‘ -4
-------
1 1—7
Missouri. The raw water was high in carbonate content (Table 2) as were
most of the waters with significant levels of uranium. But, unlike the
other water samples, the Kansas City water underwent the highest adjust-
ment in pH during treatment, from 8.1 to 9.5. This information pointed
to further investigation of the effects that raising the pH has on the
Table 2. Alkalinity versus uranium content of the municipal water
Sample caco 3 b
mg/L
u
pg/L
Midland, Mich.
8.0
82
0.27 ± 0.03,
it a
3
Houston, Texas
70
26
0.32 ± 0.14,
it a
3
Kansas City, Mo. — Raw
8.0(7.9)
172(193)
5.33 ± 0.23,
n a
3
intermediate
Final
1
2
(8.6)
(9.9)
(9.9)
(110)
(100)
4.86
4.73
4.07 ± 0.30,
a
4
Lincoln, Neb.
8.0(7.4)
153(155)
7.29 ± 0.47,
a a
3
San Diego, Calif.
Alvarado — Raw
7.6(7.7)
100(97)
1.68, 1.90
Intermediate
(8.2)
(101)
2.12
Final
(8.2)
(103)
2.31, 1.80
Escondido Vista — Raw
8.0
102
8.15 ± 0.39,
n a
3
(Calif. aque— Intermediate
(115)
6.53
duct) Final
(7.6)
(118)
6.25, 5.60
Otay —Raw
7.9(8.0)
110(122)
1.05, 1.00
Intermediate
1.22
Final
(8.2)
(121)
2.10, 2.30
Sweetwater — Raw
8.0(8.1)
135(120)
4.07, 3.30
Intermediate
(7.7)
(120)
4.25
Final
(7.8)
(120)
3.30
Los Angeles, Calif.
Hawthorne
7.9
272
0.14, 0.29,
<0.10
Jensen
0.30, 0.27
Long Beach — Raw
8.0
146(134)
0.32
Final
119
1.59 ± 0.33,
n a
3
Wey nouth
8.1
114
6.61, 6.10
Denver, Cole.
Marston
7.5
48
1.60, 1.60
Moffat — Raw
7.5
29
15.9 ± 1.58
Intermediate
1
2
7.2
7.8
20
32
3.90
4.90
Final
7.6
28
4.00
Salt Lake City, Utah
Big Cottonwood
(8.1)
(130)
0.90
Little Cottonwood
(7.9)
(136)
1.70
City Creek
(7.9)
(198)
1.00, 0.80
Phoenix, Ariz.
Val Vista
8.5
218
4.45, 4.30
Verde
8.4
206
4.30, 4.10
Squaw Peak
8.0
115
1.70
°pH recorded at tine of analysis; p H reported by sampler in parentheses.
bAlkalinity reported by sampler in parentheses.
-------
11-8
uranium levels in the sample water. Several samples were selected which
contained significant levels of uranium such that important losses could
be detected (Table 3). Four 100—mL aliquots of sample were placed in
separate glass containers and the pH in each was adjusted to 10.0, 10.5,
11.0, and 11.5, respectively, with sodium hydroxide. Depending on the
sample, as the pH was raised, a white precipitate appeared in the solu-
tion and collected on the bottom of the container. After a few minutes,
the sample solution was filtered (0.22—pm pore size) and then analyzed
by first preconcentrating the uranium in the solution onto anion exchange
resin, followed by neutron activation analysis of the resin. The results
given in Table 3 show that by raising the pH of the water, the uranium
level may be decreased, with the amount of decrease depending on the
sample.
Uranium removal was more accentuated in the sequence of samples
taken from the tiof fat Plant in Denver, Colorado. The average uranium
level for the raw water sample was .S.9 pg/L. However, the uranium level
in the postcoagulation sample was only 3.9 pg/L, a 75% decrease from the
raw water value. The pH of the raw water sample was 7.5, while the pH
of the intermediate sample was 7.2. Alkalinity was in the 20 to 30 mg/L
range (Table 2). At the point where the first intermediate sample was
taken, the raw water had been treated with aluminum sulfate, but no lime
had been added. To determine if the alum coagulation was actually
responsible for the uranium removal, an amount of alum equal to 10 mg/L
was added to a 100—a aliquot of the raw water sample, and the solution
was allowed to equilibrate with stirring for 20 minutes. Upon addition
of the alum, the pH decreased to 7. Then the solution was filtered
(0.45—pm pore size), and after additions of ammonium carbonate and
methanol, it was passed through an anion exchange column. Neutron
activation analysis of the resin showed 3.6 pg/L uranium, a 77% decrease
from the uranium level in the raw water.
The Kansas City and Denver (Moffat) results represent opposite
effects. For Kansas City, the loss of uranium, although slight, was
related to the loss of carbonate (about 50%, Table 2) during softening.
The Denver result was a consequence of a low alkalinity, low p11 sample
(this water is a mixture of softer water from the mountains and harder
-------
11—9
Table 3. Effect of raising the pit followed by filtration
(0.22 m) on the uranium concentration in water from
U.S. municipal treatment plants
Location Treatment stage pH U, ugIL
Kansas City, Mo. Intermediate 1 8.0
10.0 5.20
10.5 5.00
11.0 2.60
11.5 2.10
San Diego, Calif. Intermediate 7.7 6 • 53 a
Escondido Vista 10.0 5.20
10.5 2.40
11.0 1.40
11.5 0.80
a
Phoenix, Ariz. Intermediate 2 8.2 4.32
10.0 3.80
10.5 1.20
11.0 0.36
11.5 0.38
Lincoln, Neb. Intermediate 2 8.4
12.0 2.80
Long Beach, Calif. Intermediate 1 8.0
Weyinouth 11.0 1.40
12.0 2.10
aDta from Table 1.
water from a small reservoir receiving uranium from a mining operation)
being exposed to the aluminum hydroxide. The uranium was present in
an “unstable” state because of blending and probably would have been
slowly lost during storage even if treatment had not occurred. Thus
this was an exception to the normal stability of uranium during treat-
inent.
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11—10
2. EXAMINATION OF AVAILABLE REMOVAL METHODS
Numerous studies have been conducted concerning the recovery of
uranium from water under special circumstances, but none of the studies
attempted concerned the removal of uranium from drinking water through
existing municipal water treatment processes (see Part I of this document).
Several methods developed for uranium analysis and decontamination pro-
cesses (Hodge, 1975; Weissbuch, Cotrau, and Veliescu, 1969) could readily
become part of a water treatment procedure, but the methods should be tested
and optimized because the chemical conditions and objectives of the methods
are drastically different from those of municipal water treatment systems.
The objectives of this part of the studies were to evaluate
1. water softeners [ Ca(OH) 2 , MgCO3] as coprecipitators for uranium,
2. coagulants [ Fe 2 (SO ) 3 , FeSOi. 1 , Al 2 (SO . 4 ) 3 J as coprecipitators for
uranium,
3. adsorbents (Ti0 2 , activated charcoal) for uranium removal, and
4. ion exchange columns for uranium removal.
Interpretations of the results, particularly for the lime—softening
and coagulant experiments, are focused on the effectiveness of the chemi-
cals and optimum conditions for uranium removal. The optimum conditions
of the chemical treatments may or may not coincide with the optimum con—
ditions for the lime—softening and coagulation processes applied by
water treatment plants. The practicality of the results should, there-
fore, be examined by laboratory jar tests and pilot plant tests of waters
from different sources.
The water used in these studies was taken from Pond 3513, the former
final low—level radioactive waste settling basin of the Oak Ridge National
Laboratory. After its retirement as a waste effluent settling basin, the
pond water was subjected to a rapid turnover by rainwater runoff. We
selected this water because the uranium concentration was relatively
higher than that of any other natural bodies of water in this region.
Analytical data of the water are given in Table 4. The composition,
other than radionuclides, of the water is not greatly different from
that of the surface water in this area (ORNL—2557, 1959). The pH and
concentration of carbonate and other ligand species in the pond water
-------
‘I—il
Table 4. Chemical composition of pond (3513) watera
Constituent Concentration
(mg/L)
Total alkalinity (as CaCO 3 ) 100.5
Ca 21.4
Mg 12.9
U 0.08
K 1.9
Na 7.0
Si0 2 2.7
NO 3 0.06
so 14.1
Cl 7.8
PO (inorganic) 0.02
Dissolved oxygen 12.5
pH 8.9
1980 aAverage concentration from January to July,
suggest that the dominant uranium in the water would be uranyl carbonates
[ u0 2 (C0 3 ) 2 2 , uo 2 (c0 3 ) 3 4 ] (Langmuir, 1978). Such uranyl carbonate
species are expected to be conunon in most alkaline surface and well waters
used as raw water by municipal water treatment plants. The pond water
was collected through a filter into 25—L, polyethylene plastic jars using
a submerged water pump. The collected water was passed through Whatman
No. 42 filter paper to remove fine suspended particulates. Synthetic
waters containing various amounts of uranium and other elements were
also used in part of the experiments.
Because an alpha emitter, can be difficult and time consuming
to determine, a ganuna emitter, was used to measure the effectiveness
of each treatment in removal of uranium. The solution to be used in each
experiment was spiked with 237 U and equilibrated overnight so that the
237 j would have the same chemical form as the natural uranium in the
-------
11—12
water. Two aliquots of the solution, removed before the treatment tests,
were used as standards for the experiment. A portion of the solution was
removed before the spiking for a background (Bkgd) count. The 237 U was
measured by a Packard Auto—Ganuna System incorporating a dual thallium—
activated Nal scintillation crystal/dual photomultiplier tube detection
system. The counting error, expressed as 1001, where N is the total
number of counts, was always less than 5%. Because the 237 U has such a
short half—life (t ” 2 6.75 d), error introduced by lags in the counting
time for a set of samples was minimized by counting a standard before
and after testing each set of samples and using the average counts for
calculation of uranium removal. The following relationship was used to
determine uranium removal.
% 2 7 j removed = ( Standard count — Bkgd) — (Sample count — Bkgd ) 100
(Standard count — Bkgd)
2.1 URANIUM REMOVAL BY ALKALINE EARTH WATER SOFTENERS
Lime and lime—soda softening have been a standard practice to remove
calcium and magnesium salts, which are the main source of hardness in
water. The softening process produces a sludge composed of calcium and
magnesium hydroxide. Magnesium carbonate treatment, in addition to lime
treatment, was introduced to improve coagulation as well as to recycle
the lime and magnesium carbonate (Thompson, Singley, and Black, 1972).
Under proper conditions, removal of Ba, Ra, and other trace heavy metals
during water—softening processes has also been obtained (Sorg, Csanady,
and Logsdon, 1978; Sorg and Logsdon, 1980). Removal of uranium from
seawater by sodium hydroxide treatments was observed (Hodge, 1975), but
the removal process was not well understood.
2.1.1 Lime [ Ca(OH) 2 ] Treatment
AliquotS (200 mL) of 237 U—spiked pond water were treated with vary-
ing amounts of analytical grade Ca(OH) 2 . The lime dosages applied in
this experiment were near the range of dosages routinely used by munici—
pal plants. About 85 to 90% of the uranium in the pond water was re-
moved by the lime treatments (Table 5). The final pH of the water was
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11—13
Table 5. Removal of uranium from pond water by Ca(OH) 2 treatmenta
Ca(OH) 2
dose
(mg/L)
50
100
150
200
250
%
U
removed
86
85
87
87
90
Final pH
10.6
11.1
11.3
11.5
11.5
alnitial U concentration: 83 pgIL.
between 10.6 and 11.5. The pH of the water depended on the lime dosage,
but the differences in uranium removal were minimal. Since the effi-
ciency (85—90%) could depend on the total alkalinity, it is difficult to
conclude whether the lime treatment is sufficient to remove uranium for
all municipal water treatment plants. The concentration of indigenous
uranium in the pond water was 0.083 mg/L, and the uranium concentration
in the treated water was expected to be 0.008 mg/L, which is probably
higher than in many raw waters used by municipal treatment plants.
2.1.2 Ca(OH)2 and MgCO3 Treatment
To investigate the role of MgCO 3 addition to a Ca(OH) 2 softening
system on uranium removal, varying amounts of MgCO 3 and Ca(OH) 2 powder
were added to 200—mL portions of the 237 U—spiked pond water. The final
pH of each batch was measured after a 20—mm stirring period, and the
237 j activity remaining in the solution was measured after filtration
with 0.45—mm—pore filter.
The Ca(OH) 2 and MgCO 3 treatment results (Table 6) indicate that:
(1) At a lower dose of Ca(OH) 2 (50 mgIL), the NgCO 3 additions reduced
the effectiveness of Ca(0H) 2 for uranium removal (compare with
Table 5).
(2) At a Ca(OH) 2 dosage higher than 100 mg/L, the percent of uranium
removal increased with increase in MgCO 3 dose.
(3) At a given MgCO 3 dose, the increase of Ca(OH) 2 dose from 100 to
250 tng/L did not influence the efficiency of uranium removal.
-------
11—14
Table 6. Removal of uranium from pond water by combined
Ca(OH) 2 and MgCO 3 treatnienta
MgCO 3 dose
(mg/L)
Ca(OH) 2
dose
(mg/L)
50
100
150
200
250
10
% U removed
Final pH
32
10.6
90
11.0
90
11.1
88
11.3
89
11.3
40
% U removed
Final pH
9.0
9.8
95
10.9
95
10.8
94
11.1
94
11.4
80
% U removed
Final pH
24
10.3
98
10.8
93
10.7
98
10.9
98
11.2
120
% U removed
Final pH
15
9.9
99
10.6
99
10.8
99
11.0
99
11.2
a lnitial U concentration: 83 iig/L.
(4) The volume of floc in so1utio appeared to increase with increasing
dosage of MgCO 3 and Ca(OH) 2 .
(5) The critical pH in the MgCO 3 —Ca(OH) 2 system appeared to be near
10.6. Above pH 10.6, with sufficient dosage of MgCO 3 and Ca(OH) 2 ,
more than 98% of the uranium was removed by the floc.
(6) The optimum dosages of MgCO 3 and Ca(OH) 2 to remove uranium from the
pond water would be 120 and 150 mg/L, respectively.
The dependency of the uranium removal on MgCO 3 and Ca(OH) 2 dosage
added to the pond water suggested that water composition, particularly
indigenous Mg, Ca, and carbonate concentration, would be a factor deter-
mining the optimum dosage of the treatment chemicals.
2.1.3 pH and Uranium Concentration Effects
To investigate the effect of pH on uranium removal, solutions con-
taining 100 mgfL Ca(OH) 2 or 50 mg/L MgCO 3 were prepared from the 237 j...
spiked pond water, and then the pH of the solutions was adjusted to 9,
10, 11, and 12 with HCI and NaOH solutions. For both Ca(OH) 2 and MgCO 3
solutions, the highest uranium removal was obtained at pH 11, and the
uranium removal was reduced at pH 12 in both solutions (Table 7). The
-------
11—15
Table 7. Effect of pH at a given Ca(OH) 2 and M CO3 dose
on uranium removal from pond water
Adjusted
pH
% U removed
Ca(OH) 2
MgCO 3
100 mg/L
50 mgfL
9
24
0
10
68
0
11
80 b
59
12
48
30
U concentration: 83 pgfL.
bcompare with Table 5.
Ca(OH) 2 solution had a higher removal of uranium than the I4gC0 3 solution
at the given pHs.
In order to study the effects of variations in the indigenous uranium
concentration on the efficiency of uranium removal by Ca(OH) 2 and MgCO 3
treatments, varying amounts of 238 U were added to 200—niL aliquots of a
40—mg/L MgCO 3 solution which ha.. been spiked with 237 U. The batches
were allowed to equilibrate overnight. After equilibration, 0.02 g of
Ca(OH) 2 were added to each 200—niL batch, and the pH was measured. If it
was necessary to raise the pH to 10.9, NaOH was used. At this pH and
dosage of MgCO 3 and Ca(OH) 2 , previous results (Table 6) had shown >90%
uranium removal from the pond water. After a 20—mm stirring period,
each solution was filtered (0.45—pm pore size) and the 237 U measured.
The results shown in Table 8 indicate that at uranium levels of 2.4
pg/L or less, the efficiency of uranium removal is still 907. or better.
However, at uranium levels greater than 2.4 pg/L, the efficiency of
uranium removal decreases drastically. Section 1 of this report demon-
strates that some municipal treatment plants may process raw water with
uranium levels which fall in the 0.24— to 24—pg/L range. Inthese cases,
treatment modifications, such as higher dosage of MgCO 3 and Ca(OH) 2
(Table 6), may be necessary to maximize the efficiency of the uranium
removal.
-------
11—16
Table 8. Efficiency of uranium removal by Ca(OH) 2 and MgCO 3
treatment at varying uranium concentratiOnSa
Uranium concentration
( g/L)
2.4
x
lO
2.4
x
102
2.4
x
101
2.38
23.8
238
% U removed
97
97
93
90
58
54
aDosage: 40 mg/L MgCO 3 and 100 mg/L Ca(OH) 2 .
2.1.4 Discussion
By the addition of Ca(OH) 2 to the pond water, most of the CO 2 and
bicarbonate alkalinity is converted to calcium carbonate. These well—
known stoichioinetric reactions are:
CO 2 + Ca(OH) - CaCO 3 (s) ÷ H 2 0 (1)
Ca(HCO 3 ) 2 + Ca(OH) 2 - 2 CaCO 3 (s) + H 2 0 (2)
Magnesium bicarbonate in the water is also converted to magnesium
carbonate and then to magnesium hydroxide on further addition of Ca(OH) 2
as shown by:
Mg(HCO 3 ) 2 + Ca(OH) 2 - MgCO 3 + CaCO 3 (s) + H 2 0 (3)
MgCO 3 + Ca(OH) 2 - Mg(OH) 2 (s) + CaCO 3 (s) (4)
The addition of MgCO 3 to the system requires a higher dosage of Ca(OH) 2
to remove the carbonate as CaCO 3 and to form magnesium hydroxide, but
such additions increase the volume of the flocs composed of Ng(OH) 2 and
CaCO 3 . The gelatinous Mg(OH) 2 floc acts as a coagulant and thus improves
flocculation of the suspended materials (Thompson, Singley, and Black,
1972). Based on the jar test results, Thompson, Singley, and Black (1972)
found that good floc formation took place at a pH above 11.0.
Equations 1 through 4 suggest that as the Ca(OH) 2 dose increases,
the CO 2 and carbonates in the water become depleted and could change the
chemical state of the uranium species. The uranium in the pond water is
-------
11—17
expected to be present as uranyl tricarbonate, uo 2 (CO 3 ) 3 . The
uo 2 (CO 3 ) 3 , known as a relatively stable complex, could be converted
to uranyl hydroxide in the carbonate—depleted environment by addition
of Ca(OH) 2 . Ideally, the reaction would be:
3 U0 2 (CO 3 ) 3 + 9 Ca(0H) - (U0 2 ) 3 (0H) 5 + 9 CaCO 3 (s) + 13 011 . (5)
In the MgCO 3 —Ca(OH)2 system, the chemical reactions could be:
3 Uo 2 (C0 3 ) 3 + 10 Ca(OH) 2 + NgCO 3 - (U0 2 ) 3 (0U) 5 + + Mg(OH) 2 (s) + (6)
10 CaCO 3 (s) + 130H
Such conversion from carbonate to hydroxide was not experimentally proved
but is expected to be controlled by the ratio of the free C0 3 2 and 0H
in the water. Nevertheless, the two uranyl species are drastically
different not only in terms of ligands but also in charge and molecular
size of the species. As a result, adsorption or precipitation behavior
of the uranyl species is expected to be different.
To answer a part of the questions regarding uranium removal niech—
anisms, the following experiments were conducted with distilled—deminera
lized water instead of the pond water.
(1) Magnesium solution: 0.005 N MgC1 2 , 0.05 M NaC1, 48 ig/L 238 U, and
trace 237 U spike.
(2) Magnesium bicarbonate solution: 0.005 M MgC1 2 , 0.005 M NaHCO 3 ,
0.05 M NaCl, 48 pgIL 238 U, and trace 237 U spike.
(3) Calcium solution: 0.005 M CaC1 2 , 0.05 M NaC1, 48 vigIL 238 U, and
trace 237 U spike.
(4) Calcium bicarbonate solution: 0.005 M CaC1 2 , 0.005 M Na}1C0 3 ,
0.05 M NaCl, 48 pg/L 238 U, and trace 237 U spike.
Two—hundred milliliters of each stock solution was transferred into
a 250—mL bottle and NaOU added. The experimental results of the magnesium
and magnesium bicarbonate solutions are plotted as percent removal of
uranium versus p1-I and versus volume of 1 M NaOH added to the batches
(Figs. 1 and 2).
-------
11—18
PH OF SOLUTI ON
Al
ORNL- DWG 81 10872 ESD
£2
Fig. 1. Percent uranium (initial uranium concentration, 48 pg/L)
removed from 5 x 1o M MgC1 2 5 x 10 M MgC1 2 —5 x 10 M NaHCO 3 , and
5 x 10 M CaCl 2 —5 x M NaHCO 3 solutions at varying pHs.
More than 90% of the uranium in the magnesium solution was removed
between pH 10.3 and 10.9 or by an addition of 0.2 to 2.0 mL of 1 M NaOH
to the 200—mL magnesium solution. In the magnesium—bicarbonate solution,
the highest uranium removal was 83% at pH 10.9 with an addition of 2 mL of
1 M NaOH. The magnesium—bicarbonate solution consumed more NaOH and
required a higher pH than the magnesium solution to remove uranium. In
both solutions, the excessive NaOH or a pH higher than 11.4 reduced the
efficiency of uranium removal. A separate batch which had 0.1 M magne-
sium and bicarbonate was prepared to examine the role of hydromagnesite
[ Mg (C0 3 ) 3 (0H) 2 • 3 H 0] precipitate on uranium removal. At an adjusted
pH of 9.5, the uranium removal was only 13%.
100
Go
0
60-
-4
z
F-’ 40
z
U
rz
20
0
LEGEND
v = MG CHLORIDE
a=MG CHLORIDE—NA BICARBONATE
o=CA CHLORIDE—NA BICARBONATE
0 ’•••
- -
I ..
8
9
-------
ORNL - DWG Si - 10574 LW
£
t
0 0.5
1.5 2 2.5
VOLUME, ML OF M NAOH
Fig. 2. Percent uranium (initial uranium concentration, 48 pg/L)
removed from 5 x iO M mgCl 2 and from the mixture of 5 x 1O M
mgCl 2 —5 x ? ? aHCO 3 solution after addition of varying amounts of
1 M NaOH solution.
In the calcium solution, the uranium removal was inconsistent in the
p1! range 8.5 to 11.5, but it was less than 15%. As the pH of the calcium
bicarbonate solution was raised to above 8.7, crystalline CaCO 3 was pre-
cipitated. The uranium removal from the calcium bicarbonate solution
increased with an increase in pH, but the highest value observed was 30%
at pH 11.5 (Fig. 1).
11—19
100
0
I-
z
0
0..
80
60
40
20
0
LEGEND
v=MG CHLORIDE
= MG CHLOR DE-NA BICARBONATE
1
3
3.5 4
-------
11—20
The results from the pond water and synthesized water experiments
demonstrated the following:
(1) At the higher pH (10.7—11.3), the Mg(OH) 2 precipitate played a major
role in uranium removal. The presence of excessive carbonate lowered
the efficiency of removal. Removal of the hydroxide uranyl form
[ (U0 2 ) 3 (OH) 5 + or a monomeric analog] by the Mg(OH) 2 floc seems to
dominate over the uranyl carbonate form, U0 2 (CO 3 ) 3 , but this is
not conclusive.
(2) In the pH range between 8.5 and 10.6, the calcium and magnesium
carbonate precipitates could remove some uranium, but the efficiencies
were very low.
(3) The high removal efficiency by Ca(0II) 2 treatment without MgCO 3 in
the pond water resulted from the presence of indigenous magnesium
in the water. Therefore, the magnesium is an essential ingredient
to remove uranium from natural waters by lime treatment.
2.2 URANIUM REMOVAL BY COAGULANTS
Most of the municipal water treatment plants have been using coagu-
lation chemicals to remove turbidity, color, and organic matter from
the raw waters. The frequently used coagulation chemicals are aluminum
sulfate, Al 2 (S0 ) 3 , ferric sulfate, Fe 2 (SO ) 3 , and ferrous sulfate, FeSO 1 . 1 ,.
Hydroxide sols of these chemicals form positively or negatively charged
gelatinous flocs, and the charge and stability of the flocs are pH
dependent.
To test the effectiveness of the coagulants for uranium removal,
237 U—spiked pond water was prepared in the same manner as the lime—soft-
ening experiment. A specified dosage (0.5 to 25 mg/L) of the coagulants
was added to 200—mL aliquots of the pond water, and the pH of the solutions
adjusted to 4, 6, 8, and 10 using 1 M NaOH and HC1 in order to determine
the optimum dosage and pH for the coagulant treatments. The chemicals
used in these experiments were analytical grade. After a 20—mm stirring
period, each solution was filtered (0.45— iin pore size), and the 237 U
remaining in solution determined.
-------
I I— 21
2.2.1 Fe 2 (SOi )3 Treatment
As expected, the removal efficiency by ferric sulfate depended upon
both the dosage and the equilibration pH of the solution (Table 9 and
Fig. 3). The removal percentage increased with increase in dosage, and
this trend was more evident in the solutions which had a pH of 6 or 10.
There were no considerable differences in uranium removal among the
dosages above 10 ing/L at pH 10. At pH 6 and 10, the uranium removal
efficiency in the batches containing 25 mg/L ferric sulfate was about
88%, but it was only 40 and 20% at pH 8 and 4, respectively.
2.2.2 FeSOL+ Treatment
The experimental results of the ferrous sulfate treatments (Table 9
and Fig. 4) were similar to those of the ferric sulfate treatment in terms
Table 9. Percent uranium removal by Fe 2 (S0 1 ) 3 , FeSO , and Al 2 (S0 ) 3
coagulants with varying
Initial
pH
Dosage
(mg/L)
U
removed
(%)
Final pH
Fe 2 (S0 4 ) 3
FeS0
A1 2 (S0 14 ) 3
Fe 2 (S0 ) 3
FeSO 1
A1 2 (S0 1 j 3
4
0.5
5
10
15
20
25
7
14
8
13
17
18
6
8
11
21
26
33
7
9
6
15
2].
21
4.1
4.3
3.8
4.0
4.1
4.0
4.2
4.2
4.1
4.2
4.2
4.1
4.4
4.2
4.4
4.7
4.8
4.8
6
0.5
5
10
15
20
25
16
43
63
76
84
89
14
24
33
42
52
44
7
30
51
69
80
88
6.2
6.4
6.2
6.2
6.1
- 6.2
6.1
6.0
6.1
6.2
6.2
6.2
6.2
6.2
6.1
6.1
6.].
6.2
8
0.5
5
10
15
20
25
1
4
17
21
33
43
6
7
12
11
15
20
0
2
9
17
25
48
8.4
8.2
7.9
8.0
7.9
7.8
8.1
8.1
8.1
8.1
8.0
8.0
8.0
7.9
7.9
7.9
7.9
7.8
10
0.5
5
10
15
20
25
1
27
83
86
80
87
2
32
57
84
92
93
8
71
95
98
98
96
10.0
10.0
9.9
10.0
9.5
10.0
10.1
10.0
10.0
10.0
9.9
9.9
10.0
9.9
9.8
9.7
9.7
9.7
aInltial U concentration: 83 gIL.
-------
11—22
0 f4L . DWG 81 10878 SU
100
pH
‘=6
0 •0.... -
$0 o=1O
0
-
e0,
z
V
40
z
0
‘I
I,
‘I
20
0
FERRIC SULFATE DOSE, MG/L
Fig. 3. Percent uranium removed from pond water (initial uranium
concentration, 83 i .ig/L) as a function of doses (rnglL) of Fe 2 (SO ) 3 and pH.
of the dependency of uranium removal on pH and dosage. At pH 10, 20, and
25 mg/L, dosages of ferrous sulfate removed more than 90% of the uranium
in the solutions. The coagulant was less effective at pH 6, but at pH 6
the removal was higher than at pH 4 and 8.
2.2.3 A12(S0L+)3 Treatment
The uranium removal efficiency at pH 10 was 95% with 10 mg/L or higher
dosages of aluminum sulfate (Table 9 and Fig. 5). At pH 6 removal increased
with increase in dosage, and 87% removal was obtained with a 25 mg/L
dosage. On the other hand, only 48% and 21% of removal was obtained at
pH 8 and 4, respectively, with the same dosage. These results can be
compared to those of 10 mg/L alum addition to the process of the Mof fat
-------
100
00
0
00
z
40
z
0
20
0
J. £ . J
O PdL.DWG S1-1O$77 ESO
Fig. 4. Percent uranium removed from pond water (initial uranium
concentration, 83 iig/L) as a function of doses (mg/L) of FeSOt 4 and pH.
plant (Denver) discussed in Sect. 1. Under lower alkalinity conditions
(Table 5) and near neutral pH, about 75% of the uranium was removed.
2.2.4 Discussion
The physiochetnical properties of the hydroxides [ Fe(OH) 3 , Fe(OH) 2 ,
Al(OH) 3 ] formed from ferric, ferrous, and aluminum sulfate during the
coagulation process are well known. The stability, solubility, and
reactivity (adsorption) of the hydroxides are pH dependent. The charge
characteristics (zero point of charge of the hydroxides) are different,
but the hydroxides have positive charge in the acid range, mixed (neutral)
charge at pH 5—7, and negative charge in the alkaline range in general.
FERROUS SULFATE DOSE, MG/L
-------
11—24
ORNL DWG 81 10876 ESD
Fig. 5. Percent uranium removed from pond water (Initial uranium
concentration, 83 iig/L) as a function of doses (mg/L) of A1 2 (S0 ) 3 and pH.
The pH dependency of the distribution of uranyl species in natural water
Is also well known.
Positively charged UO would dominate at pH 4, neutral uranyl
carbonate UO 2 CO at pH 6, and negatively charged uranyl carbonates,
UO 2 (CO 3 ) and U0 2 (CO 3 ) , at pH 8 and above. At pH 10, the Uo 2 (Co 3 )
species Is known to be stable, but (U0 2 ) 3 (OH) 5 + would be a dominant
species in carbonate—depleted water. The carbonate in pond water would
be depleted by CaCO 3 precipitation during the coagulant treatment process.
3 UO 2 (CO 3 ) + 15 CaZ+ + (Fe, Al) 2 (SO ) 3 + 9 OH
0
z
40
z
U
10 15
ALUMINUM SULFATE DOSE, MG/L
(UO 2 ) 3 (0H + 12 CaCO 3 (s) + (Fe, Al)(OH) + 3 CaSO 4
-------
11—25
The pH dependence of the experimental results (Table 9) in uranium
removal appears to relate to the uranium species and charges of the
coagulant flocs at a given pH. Minimum uranium removal was observed
when the charge of the uranyl species was the same as the charge of the
flocs, and maximum removal occurred when the charges were opposite or
neutral (Table 10).
Although the best results of uranium removal were obtained at the
equilibration pH of 10, in practice, the stability of the flocs should
be considered. The iron hydroxide is stable in a relatively wide range
of pH, but the amphoteric aluminum hydroxide floe is unstable and dis-
solves at high pH. Therefore, unless coagulant treatment is incorporated
with lime treatment, the best results will be obtained at pH 6. A higher
dosage (>25 ing/L) of coagulartts was required to remove uranium at pH 6
than at pH 10.
2.3 URANIUM REMOVAL BY ADSORBENTS
Many adsorbents such as peat, coal, and hydrous metal oxides have
been studied as to their potential for removing uranium from seawater
(see Part I of this document). Two of these adsorbents were investigated
in this study for removal of uranium from municipal water. Titanium
oxide had already been selected for use in a pilot plant to recover
uranium from seawater (Technology Alewsletter, 1980), so it seemed a
Table 10. Suspected uranyl species and charge characteristics of iron and
aluminum hydroxide flocs at given pHs of pond water
Adjus
ted p11
4
6
9
10
Uranyl species
U0 2 2
U0 2 CO
U0 2 (C0 3 )r
(u0 2 ) 3 (OH)t
Charges of flocs
(Fe,A1)(OH)
(FeAl)(OH)
(Fe,A1)(OH) .y
(Fe,A1)(OH) ÷z
Uranium removal (%)Q Low (30)
High (88)
Low (48)
High (87)
25 mg/L dosage.
-------
11—26
likely candidate to study for use in water treatment plants. The adsorp-
tion capacity of Ti0 2 for uranium from seawater had been studied pre-
viously (Ozawa et al., 1979). Activated charcoal is employed as a taste
and odor control in many water purification plants, so it was also
investigated for removal of uranium. Its usefulness as an adsorbent for
uranium had already been demonstrated in neutron activation determina-
tions (Kuleff and Kostadinov, 1978).
Stock solutions of 10—2 and l0-’ M sodium bicarbonate and sodium
carbonate were prepared with analytical grade chemicals. The batches
containing 200 mL of the stock solution were spiked with 237 U, and the
pH of the solutions was adjusted with 1 M NaOH or HC1 solutions. Then,
each solution was added to a 250—mL bottle containing one gram of Ti0 2
(Baker Chemical Co., Ultrex) or activated charcoal (6—14 mesh, Fisher
Scientific Co.). After an overnight equilibration with shaking, the
final pH was recorded, each solution was filtered (1.2—iim pore size),
and the 237 U remaining in the solut 4 ons was determined. The free C0
concentration in the solutions was calculated using the final pH in the
following relationship:
(COt) = 4.7 x l0h1(HC0 3 /(}i’)
2.3.1 Titanium Oxide (Ti0 2 )
The uranium adsorption on titanium oxide was influenced by both
pH and C0 concentration (Table 11). In the solution with a free C0
concentration of 3 x l0 M or less, 96% or more of the uranium was
adsorbed on the titanium oxide. The pH of the solutions ranged from
5.1 to 6.9. As the free CO concentration and pH increased, the percent
uranium adsorption decreased. At a similar pH (8.3—8.5), as the free
C0 concentration increased from 9 x 10—6 M to 1 x l0 M and 1 x l0
M, the uranium adsorption decreased from 93 to 85 and <1%, respectively.
2.3.2 Activated Charcoal
The results (Table 12) indicated that as pH and CO concentration
in solution increased, the uranium adsorbed on the activated charcoal
decreased. At pH 6.4 or lower and free CO concentration less than
-------
11—27
Table 11. Effects of carbonate concentration and pH
on the adsorption of uranium on titanium oxidea
Initial
carbonate
concentration
(M)
Ca1culat d
free CO 3 —
concentration
(l4)2
Final pH
237 U
adsorbed
(%) —
0.001 NaHCO 3
3 x
6.9
96
9 x 10—6
8.3
93
2 x
9.7
61
0.01 NaHCO 3
6 x 10_B
5.1
98
4 x 10—6
7.0
67
4 x lO
8.9
<1
0.001 Na 2 CO 3
2 x 10—8
5.5
98
1 x i0
8.3
85
2 x lo-
9.8
46
0.01 Na 2 CO 3
6 x 10_8
5.1
98
1 x
8.5
<1
6 x lo
10.5
<1
aSOlld to solution ratio: 200 niL solution per 1 g
Ti0 2 (Baker Chemical Co.). Initial uranium concentration:
trace 237 U ( 2 x 10.6 pg/L). Equilibration time: l6 h
with shaking.
bCalculated from measured final pH..
106 M, more than 95% of the spiked 237 U was removed from solution by
the activated charcoal. On the other hand, at pH 9.1 and 6 x 10 ’ M
C0 concentration, the adsorbed uranium was less than 1%. At approxi-
mately the same pH (9.1—9.3), the adsorbed uranium was about 70% in a
solution containing 6 x i0 M of C0 .
-------
11—28
Table 12. Effects of carbonate concentration and pH on
the adsorption of uranium on activated charcoala
Initial
carbonate
concentration
(M)
Calculated
free C0 3 2
concentration
(M)b
Final pH
237w
adsorbed
(%)
0.001 NaHCO 3
7 x 10
7.6
91
9 x io
9.3
75
6 x 1o
10.1
39
0.01 NaHCO 3
3 x i0
5.8
97
1 x 1o
7.4
75
6 x 1o
9.1
<1
0.001 Na 2 CO 3
2 x 10—6
7.7
92
6 x i0
9.1
71
5 x 1o
10.4
33
0.1 Na 2 CO 3
1 x 10—6
6.4
96
1 x i —
7.4
62
6 x
9.1
<1
aSOlId to solution ratio: 200 mL solution per 1 g
activated charcoal (16-14 mesh, Fisher Scientific Co.).
Initial uranium concentration: trace 237 U (2 x 106 pg/L).
Equilibration time: 16 h with shaking.
bCalculated from measured final pH.
2.3.3 Discussion
Since most of the raw waters of concern for uranium removal which
are received by municipal treatment plants have a pH in the range of
7 to 8, uranium removal by adsorption onto titanium oxide or activated
charcoal may be possible if the carbonate concentration is low. There-
fore, the sequence of steps in the water purification process becomes
very important. Lime additions or other treatments that would affect the
-------
11—29
pH and carbonate concentration of the water would also affect uranium
removal by adsorption onto titanium oxide or activated charcoal. These
materials cannot be considered universal adsorbents. In waters where
uranium may be a problem, sorption is likely to be poor (carbonate concen-
trations may be high).
2.4 URANIIJN REMOVAL BY ION EXCHANGERS
Ion exchange resins, particularly anion exchange resins, have been
used to recover uranium from uranium mine waters (Ross and George, 1971).
The technology has proved to be the most effective and economical recovery
method, but most of the existing public water treatment systems do not
have such facilities. The application of exchange resins to remove
uranium from drinking water would, therefore, be limited to small
communities where households treat the raw water through commercially
available filtering systems.
2.4.1 Batch Test
To investigate the adsorption of uranium by anion exchange resin,
200—mL batches of the 237 U—spiked pond water were adjusted to pH 5, 7,
and 9 and then transferred to a 250—mL bottle containing a known
amount of resin (50—2000 mg, wet resin weight). After an overnight
equilibration with shaking, the solutions were filtered (1.2—urn pore
size), and the 237 U remaining in each solution was determined.
In most of the batches, 95% of the uranium in the pond water was
adsorbed by the anion exchange resin (Table 13). The adsorption of
the uranium did not depend on pH or on the carbonate concentration. Low
adsorption in the batches with 16 ing of resin did not mean that the
exchange sites were saturated with respect to the uranyl species, but
that the amount of resin was probably too small to contact all of the
available uranyl species in the solution.
To determine the effects of CaS0i on uranium adsorption by anion
exchange resin, varying amounts of resin were added to 237 U—spiked
0.001 and 0.01 H CaSOi. 1 solutions in the same manner as the pond water.
-------
11—30
Table 13. Adsorption of uranium by anion exchange
resin from waste pond water at varying pH
(Dowex 1—X2, 50—100 mesh, chloride)
Dry resin
weight
(mg)
Calculated
free CO
concentration
(M)
Final
pH
237 U
adsorbed
(%)
16
1 x 10—8
5.2
88
165
1 x 10—8
5.2
89
660
3 x 10
5.6
95
16
9 x 1O
7.1
66
165
6 x io
6.9
96
660
3 x i0
6.6
96
16
1 x 1o
9.2
68
165
1 x 1o
8.2
96
660
1 x 10
7.2
96
of solution: 200 mL. Initial uranium
concentration: 83 iig/L. Equilibration time: l6 h
with shaking.
bMi content of resin determined at 105°C:
67%.
CCalculated from total carbonate of pond water
(1.49 x M) and final pH.
The results (Tables 14 and 15) indicate that 98% of the 237 U was adsorbed
on the resin. Variations in pH and CaS0 concentration did not influence
the efficiency of uranium adsorption by the resin. The results obtained
from the batches with 16 mg of resin were inconsistent, probably due to
the extremely small solid to solution ratio, which permitted only limited
contact of the uranyl ion by the exchange resin.
Similar experiments were conducted to determine the effects of NaC1
concentration at different pHs. The results were essentially the same
as those for the CaSOi experiments (Table 16).
-------
11—31
Table 14. Adsorption of uranium by anion exchange resin
from 0.01 M CaS0 4 solution at varying pHa
(Dowex l—X2, 50—100 mesh, chloride)
Dry resin
weightb
(mg)
Final pH
237 U
adsorbed
(%)
16
5.2
99+
165
5.8
99+
660
6.8
99+
16
6.7
68
165
7.0
99
660
7.3
99+
16
8.4
98
165
8.4
91
660
8.3
98
of solution: 200 mL. Initial uranium
concentration: trace 237 U (=2 x pg/L). Equili-
bration time: =16 h with shaking.
bM. content of resin determined at 105°C:
67%.
2.4.2 Column Test
2.4.2.1 Cation Exchange Column
+ + 2+
Cation exchange columns in the H , Na , and Ca forms were prepared
using Dowex 50—X8, 50—100 mesh (exchange capacity 1.7 meq/rnL). The
columns contained 5.5 mL resin, 1 cm in diameter and 7 cm in length, and
the flow rate was 4 niL per 0.8 cm 2 per mm. A total of 2400 niL of 237 U—
spiked pond water was pumped into the H+_form column, and each 200 niL of
the effluent was collected. The analysis of the effluent batches showed
a gradual increase from 3% at the start to 7% of the total 237 U activity
in the last batch of the effluent. The pH of the effluent was 3.5 and
did not change during the elution experiment. The results suggested
-------
11—32
Table 15. Adsorption of uranium by anion exchange resin
from 0.001 N CaS0 , solution at varying pHa
(Dowex 1—X2, 50—100 mesh, chloride)
Dry resin
weightb
(mg)
Final pH
237 U
adsorbed
(¼)
16
52
93
165
6.1
98
660
6.4
98
16
6.6
99
165
6.8
99+
660
6.8
99+
16
8.8
82
165
8.5
98
660
7.6
99+
a
Volume of solution: 200 mL. Initial uranium
concentration: trace 237 U ( 2 x i0 i.ig/L). Equili-
bration time: 16 h with shdking.
bMj content of resin determined at 105°C:
67%.
that the uranyl carbonates in the pond water were changed to uranyl
cations in the acid resin bed. Removal of uranium by the Na and Ca
forms appeared to be low and could be caused by the low selectivity of
uor over Ca2+, Mg2+, and/or by slower dissociation of carbonate from
its uranyl complex.
To test this, the pH of the pond water was adjusted to 8.2, 7.0,
5.6, and 4.0 and each solution pumped, in the order of high pH to low
pH, into the Ca and Na columns. Neither column removed uranium at pH 8.2
(Table 17). At pH 7 the Ca column did not remove uranium, but the Na
column removed about 85%. The Ca column started to remove uranium from
the solution at pH 5.6; at pH 4.0, about 60% of uranium in the solution
was removed. The Na column continuously removed 70% at solution pUs of
5.6 and 4.0. The pH dependency of the removal process illustrated the
-------
11—33
Table 16. Adsorption of uranium by anion exchange resin
from 0.02 N NaC1 solution at varying pHa
(Dowex 1—X2, 50—100 mesh, chloride)
Dry resin
weightb
(mg)
Final pH
237 U
adsorbed
(%)
16
5.2
85
165
6.0
93
660
6.8
99
16
6.7
84
165
6.9
98
660
7.4
99+
16
7.7
93
165
7.8
99+
660
7.8
99+
Volume of solution: 200 inL. Initial uranium
concentration: trace 237 U ( 2 X 10 pg/L). Equili-
bration time: 16 h with shaking.
bMi content of resin determined at 105°C:
67%.
importance of uranium species in the feed solution on the removal. The
difference of removal efficiency between Ca and Na columns suggested that
the selectivity order of the cation exchange resin would be uor > Ca2+
> Na+ at acid pHs. However, the selectivity of the cation exchange resin
2+ 2 2+ -
f or DO 2 over Ca , Mg , etc., is probably not large enough to process
drinking waters.
2.4.2.2 Elution of Uranium from Cation Exchange Resin
Small plastic columns, 1 cm in diameter and 10 cm in length, con-
taining 7 mL of H+_, Na+_, or Ca2 4 _saturated cation exchange resin
(Dowex 50-X4, 50-100 mesh) were spiked with small amounts of 237 U
solution. The 237 U in the resin columns was eluted with solutions
-------
11—34
Table 17. Removal of uranium by Ca— and Na—resin column
at varying solution pHa
Influent
pH
b
Batch
number
Ca—resin
column
Na—resin
column
%
Removed
Effluent
pH
%
Removed
Effluent
pH
1
5
8.2
10
8.4
8.2
2
3
8.2
6
8.4
3
2
8.3
5
8.4
4
5
7.2
87
7.3
7.0
5
2
7.2
86
7.2
6
1
7.2
83
7.2
7
60
6.1
77
6.3
5.6
8
40
6.0
76
6.1
9
28
6.0
75
6.1
10
65
5.1
73
5.7
4.0
11
65
4.3
74
5.4
12
63
4.2
74
5.2
alnitial U concentration: 83 pg/L.
bEach batch has 100—mL effluent, and the number was the order of
collection.
containing NaC1 and (NHL) 2 C03 (Table 18). The results indicated that
1 M NaC1 plus 0.1 M (NH 1 ) 2 C0 3 solution was more effective than 5 M NaC1
solutions with or without 0.1 M (NH ) 2 C0 3 for 237 U elution from H+_ and
Na+_saturated resin columns. The lower recovery of 237 U from Ca2+_
saturated resin column could be caused by physical hindrance of calcium
carbonate precipitates on the resin surface. It is known that sorbed
uranium on cation resin can be quantitatively recovered with 2 to 4 M
HC1 or HNO 3 solution, but the recovery rate is decreased as the acid
concentration decreases because of lower selectivity of the resin for H 4
over U0 2
-------
11—35
Table 18. Percent uranium recovered from cation resin columns
by various eluting 0 i ti 0 5 a
Initial form of
U—spiked resin
Eluting solution
Column
volume 1 ’
4
8
12
16
H
1 M NaCl + 0.1 H (NH ) 2 C0 3
5 M NaC1
4.5
33.3
69.7
48.7
88.9
50.9
94.0
52.5
Ca2
1 H NaC1 + 0.1 M (NHk)2C03
5 M NaC1 + 0.1 M (N}I ,) 2 C0 3
31.1
38.1
33.6
46.6
44.3
56.7
53.0
60.9
Na
1 H NaC1 + 0.1 H (NH ) 2 C0 3
69.6
86.3
92.8
96.4
5 H NaC1 + 0.1 M (NHt ) 2 C0 3
51.7
68.6
78.6
82.8
U concentration: trace
bii volume: 7 mL.
In practice, a two—stage column system should be considered to
remove uranium from drinking water. In the first stage, one could remove
normal hardness in water using a Na+_saturated resin column. Then uor
in the effluent could be removed after dissociation of complexing anions
using a H+_saturated resin column in the second stage. The used resin
columns could be regenerated by NaC1 treatment for the Ca2+ and Mg2+
saturated resin and acid treatment after the NaC1—(NHz 4 ) 2 C0 3 elution of
2+
the U0 2 —saturated resin. One advantage of this type of dual column is
that radium would also be removed. Indeed, such a column would be pre-
ferred for the radium plus uranium case. More research needs to be done,
however.
2.4.2.3 Anion Exchange Column
An anion exchange column, 1 cm in diameter and 7 cm in length with
5.5—mL column volume, was prepared with carbonate saturated resin (Dowex
l—X2, 50—100 mesh, 0.7 meq/niL resin bed). Fifty liters of the pond
water were pumped into the column at a flow rate of 4 niL per 0.8 cm 2 per
mm. Periodically, 500—mL fractions of the effluent were collected,
and uranium alpha activity in the selected effluent batches was determined.
-------
11—36
After passing 50 L of the pond water through the column, the uranium
remaining in the effluent was less than 1% of the feed solution and did
not change throughout the experiment. The results suggested that a much
larger quantity of pond water, which would be difficult to handle in the
laboratory, would be required to completely load the resin with uranium.
The total uranium retained in the resin column was 4.1 mg
(0.083 mg/L U x 50 L x 0.99). Total exchange capacity of the 5.5 mL
anion exchange resin was 3.8 ineq. Assuming that the dominant uranium
species in the pond water is UO 2 (C0 3 ) , 5 (50% dicarbonato, 507, tricar—
bonate complex), the milliequivalent weight of the uranium is 79 mg
(238 3). Therefore, the 5.5 inL resin in the column could adsorb about
300 mg of uranium. These estimates suggest that the column treated with
50 L of the pond water had 1.4% of the resin’s capacity saturated with
uranium.
To investigate the effect of column dimension and flow rate on
uranium removal efficiency by the anion exchange column, two columns
(9 and 3 cm in length and 1 cm in diameter) were tested with varying
flow rates of 237 U—spiked pond water (Table 19). In all cases, more
than 99% of the uranium in the pond water was removed by the carbonate—
saturated anion exchange resin.
2.4.2.4 Uranium Loading Capacity of Anion Exchange Resin
The apparent failure to evaluate maximum uranium loading with the
pond water prompted another attempt with a smaller column and a higher
Table 19. Uranium adsorption by two different anion
exchange columns with varying flow rate of
237 U—spiked pond watera
Column
length
(cm)
Flow rate, mL per
0.785 c m 2 per
mm
1
3
6
12
3
99
99
99
99
9
99
99
99
99
auranium concentration: 83 pg/L.
-------
11—37
uranium concentration in the feed solution. The column was 3.6 cm long
and 0.4 cm in diameter with 0.5 mL resin volume. The 237 U—spiked feed
solution had 23.8 jg/mL (10 Al) 238 U, 10—2 N (NHi) 2 C0 3 ; the flow rate
was 5 mL per 0.1 cm 2 per mm. The calculated loading capacity of the
0.5 uiL resin (Dowex l—X8) was 48 mg of uranium if the uranium in the
solution was U0 2 (C0 3 ) . 5 form.
A gradual increase in the uranium concentration in the effluent
solution was observed (Table 20 and Fig. 6). The uranium in the effluent
was 1% of the initial concentration up to when 400 inL of the feed solution
had passed through the column. At this point, 20% of the calculated
capacity of the resin was saturated with uranyl carbonate. When the
column processed 1300 niL of the feed solution, the effluent contained
10% of the initial concentration, and 65% of the resin capacity was satu-
rated. When 1800 mL of the feed solution passed through the column, 81%
of the resin was saturated with uranyl carbonate, and 28% of the uranium
in the feed solution was in the effluent solution (Table 21). The experi-
mental data curve (Fig. 6) appears to fit an exponential curve, y = be ,
where y = percent uranium remaining in effluent, b = 0.79 as the intercept
on y, and m = 0.002 as the slope with correlation coefficient Cr) = 0.99.
Table 20. Removal of uranium by a 0.5—niL anion resin columna
Volume
effluen
(mL)
of
t
Uranium Volume
in effluent effluen
(%) (mL)
of
t
Uranium
in effluent
(%)
100
1.0
1000
5.7
200
1.0
1100
6.7
300
1.0
1200
8.3
400
1.0
1300
9.7
500
2.2
1400
12.5
600
2.6
1500
15.6
700
3.1
1600
18.1
800
3.8
1700
22.9
900
4.5
1800
28.5
a
Feed solution had 23.8 mg/L (10 N) of uranium and
10.2 M (NH ) 2 CO 3 .
-------
z
z
—
z
z
0
40
20
0
11—38
ORNL - OWG 8 - 10875 ESD
200 400 600 800 1000 1200 1400 1600 1800 2000
COLUMN VOLUMES OF EFFLUENT SOLUTION
Fig. 6. Percent uranium (initial uranium concentration, 23.8 mg/L)
passed through an anion exchange column versus cumulated column volumes
of influent solution.
2.4.2.5 Elution Experiment of Anion Exchange Resin
The same column used for the loading experiment (Table 20) was
eluted with 1.0 Al NaC1—O.5 N NaHCO 3 solution. This column contained
39 mg of uranium (Table 21); about 28% of the uranium in the column was
eluted with four column volumes (2 mL) of the solution, 60% with 14
column volumes, and 85% with 40 column volumes of the eluting solution
(Fig. 7). The remaining 15% remained with resin.
In a separate experiment, uranyl carbonate—treated anion exchange
resin columns were eluted with solutions having various compositions
(Table 22). The solutions containing 1.0 or 1.5 M NaC1 and dilute
carbonate had the highest efficiency.
I
0
-------
11—39
Table 21. Percent of uranium loading on resin column at
the selected uranium concentration in the effluent
Volume
effluen
(mL)
of
t
U in
effluent
(%)
U in
resin column
(mg)
Loading of
resin column
(%)
400
1.0
10
20
1300
9.7
31
65
1800
28.5
39
81
C-
I
z
100
00
80
40
20
0
ORNL DWG 81 10873 ESD
Fig. 7. Percent of uranium eluted from uranium—containing anion
exchange column (initial uranium loading, 39 tog) by 1.0 M NaC1— 0.5 M
NaHCO 3 solution versus cumulated column volumes of elution solution.
4 8 12 16 20 24 28 32 36
COLUMN VOLUMES OF INFLUNT SOLUTION
-------
11—40
Table 22. Percent 237 U recovered from anion resin columns
by various eluting solutions
Eluting
solution
Column
volumea
1.5
3
4.5
6
1.5
M
NaC1
+
0.1 M (NHj ) 2 C0 3
57.8
82.2
90.8
94.8
1.5
M
NaC1
+
0.5 M Na}1C0 3
52.6
76.0
83.7
88.9
1.5
M
NaC1
+
0.1 M NaS0
49.1
72.4
81.7
85.5
1.5
M
NaC1
+
0.5 M NaS0
49.3
73.1
81.5
85.2
1.0
M
NaC1
+
0.5 M NaHCO 3
50.0
75.3
90.2
95.1
1.0
M
NaC1
+
0.1 M (NHL ) 2 C0 3
42.1
74.4
84.7
90.1
5.0
M
NaC1
+
0.1 M (NH ,) 2 C0 3
35.8
49.7
57.6
64.6
5.0
M
NaC1
+
0.01 M (NH ) 2 CO 3
37.4
52.6
62.2
68.7
a 1 column volume = 7 tnL.
2.4.3 Discussion
Both batch and column results confirmed that the strong base anion
exchange resin has a very large adsorption capacity and selectivity for
uranyl carbonates, which are common chemical species of uranium in surface
and groundwaters. The uranyl carbonate adsorption on the resin column
depends on flow rate, pH, and concentration of uranyl carbonate and other
competing anions. At relatively high uranium concentrations M)
of influent solutions, loading capacity of a resin column at a point of
5% leakage in effluent decreased with increasing flow rate and carbonate
concentration (r ,l01 M) (Shankar, Bhatnagar, and Murthy, 1956). The
study also found that the loading capacity decreased with decreasing
uranium concentration, indicating that a bigger column volume would be
required to obtain the same percent leakage in the effluent from a
diluted influent as from a more concentrated one. In the presence of
vanadium in the influent solution, uranium loading capacity of anion
resin would decrease because vanadium would be held much more strongly
by the resin (Grinstead, Ellis, and Olson, 1956). Using a relatively
large anion resin column with flow rate 100 L per 1000 cm 2 per mm, Ross
and George (1971) were able to remove 98% of uranium from mining solution
-------
11—41
containing 4 x io M uranium. Total loading of uranium at the 2%
leakage point was 50% of the theoretical resin capacity.
The loading experiment of this study (Table 21) with l0 N uranium
and 10 -2 H (NH 1 3 2 C0 3 solution and a flow rate of 5 mL 10.1 cm 2 /min showed
that at 20% loading the uranium leakage was 1%, and at 65% loading the
leakage was 10% of the uranium concentration in the influent solution.
An extrapolation of the above results, obtained from relatively
concentrated uranium systems, to an extremely diluted system such as the
pond water or surface water and groundwater may or may not be valid, but
there are no other ways to determine the loading capacity and leakage
concentrat ion for the diluted system without site—specific field experi-
ments. The pond water had 3 x l0 N uranium, 10 N HCOj, and clO ’ H
of other anionic species (Table 4). The raw waters used for the municipal
water plant have l0 to l0 N of uranium. In general, the molar ratio
of uranium to total anionic species in natural waters is 1:lO , and most
of the laboratory loading (literature) studies were conducted with a
ratio between 1:102 and l:l0 . The low uranium concentration and high
anion concentration in natural waters could reduce the efficiency of
uranium removal by anion exchange column.
For estimation purpose, a very conservative value (10% maximum loading
with 1% uranium leakage) was selected. A commercial demineralizer
(Cole—Parmer, Cat. No. C—1503, flow rate, 2 gal/Mn) could hold 4400 mL
(5 lb) of anion exchange resin. The 10% loading capacity for U0 2 (C0 3 ) 2 5
of the demirieralizer is equivalent to 21.12 g of uranium. The deminer-
alizer could, therefore, process 2.6 x l0 L of the pond water (0.08 mg/L
of uranium) with <1% leakage equivalent to <0.8 vigIL of uranium in the
effluent. If a water has 0.01 mg/L, the demineralizer would process
2 x l0 L of the water and the uranium concentration of effluent would
be 0.1 pg/L. The list price of the demineralizer was $108, and the
price of replaceable anion exchange resin was $60 per 5 lb (no endorse-
ment of the Cole—Parmer unit is intended). The ganuna dose rate at one
meter for 234 Th, 234 Pa, and 23L4U can be calculated assuming a point
source. Thus 1 g of 238 U (in equilibrium with 23413 in the water) on an
anion column would result in an annual dose of 0.02 mrem (natural back-
ground is near 100 to 200 mrem/year).
-------
11—42
The recovery rate of uranium from anion resin columns by eluting
solutions was relatively low; less than 60% was eluted with 1.5—column
volumes of the solutions. Under the best conditions, 90% was recovered
with elutions of 4.5 column volumes. The elution efficiency could be
improved by increasing retention time of the solution on the resin. Ross
and George (1971) were able to recover 98% uranium from 50% saturated
resin using 1.9 column volumes of 1.1 M NaC1—O.O5 M NaHCO 3 solution at
a resin retention time of 3.5 h. Since the elution efficiency would also
depend on the column loading and if the anion exchange process would be
considered as one of the options for municipal water treatment, further
detailed studies should be conducted. For the household—type demineral—
izer column, recycling the resin by users may not be practical.
2.5 URANIUM REMOVAL BY REVERSE OSMOSIS
Reverse osmosis (RO) water treatment systems are generally effective
for the removal of most inorganic species in feed waters (Sorg and Logsdon
1980). A study of RO performance in the separation of uranium from
synthetic mine solutions indicated that the RO system with cellulose
acetate membranes was highly effective (above 96% rejection rate) in the
uranium sulfate concentration range of 100 to 800 mg/L in feed solutions
(Sastri and Ashbrook, 1976). The performance of the RO system would,
however, depend on membrane material, system construction, operation
condition, feed solution composition, etc.
In this experiment, the uranium removal efficiency of an RO system
was examined using a commercial cellulose acetate membrane module (2—in.
deep by 24—in, long, membrane area 11 ft , Cole—Parmer Instrument Co.,
Cat No. C—1501—90). A cylindrical plastic compartment (2.3 in. deep by
24 in. long) for the module and feed solution was fabricated. The flow
rate of effluent through the membrane was 0.3 gal/h and was regulated by
N 2 gas pressure (20 psi).
About 4 L of prefiltered (Whatman No. 42) pond water containing 83
pg/L uranium passed the RO membrane unit. Then 1.5 L of 237 U—spiked
pond water was put through the unit. Five 200—mL aliquots were collected
successively after discarding the first 500 mL of the effluent of- the
-------
11—43
spiked pond water. The results (percent removal) of spiked 237 U and
indigenous 238 U from the feed solution are given in Table 23.
The percent uranium removal (rejection rate by the RO unit) decreased
as the filtration progressed. Although the removal of uranium was high
(>90%) for the first 4 L (data not shown but inferred from initial col-
lection data, Table 23), apparently only about 5 L can be treated at close
to 907. removal with this design. This is not as satisfactory as the
anion exchange resin column. Since a principal objective of this test
was to evaluate the potential of the method, the capacity is not a factor,
and it can be concluded that RO can remove uranium to below 90% of influent
concentrations and, therefore, is a viable option for treatment purposes.
Table 23. Percent of uranium removed by a reverse osmosis module
Fraction of effluent
collected (mL)a
237 U
removed (%)
238 U
removed (%)
500—700
93
94
700—900
87
91
900-1100
87
85
1100—1300
85
85
1300—1500
79
86
aAfl initial 4 L was passed prior to the initiation of sample
collection after the introduction of 237 U into the source water.
-------
11—44
3. SUMMARY OF RESULTS
The total dissolved uranium concentration in the water samples
provided by 20 public water supplies was in a range from 0.1 to 9.1 ig/L.
The pHs of the raw and treated waters were between 7.4 and 9.5. The
lack of noticeable differences in uranium content among raw water, inter-
mediate stage, and product water samples indicated that municipal water
treatment practices are not effective in removing uranium. Supplementary
work suggested that the treatment conditions established for the reduction
of turbidity and hardness by municipalities are not optimum for the
removal of uranium.
Batch tests of water softener, Ca(OH) 2 , showed that lime treatment
alone could reduce uranium concentrations of the tested water (83 ig/L
uranium) by 85 to 90% at pH between 10.6 and 11.5. The removal effi-
ciency was improved up to 99% by the addition of MgCO 3 to the water.
Uranium was removed from solution y coprecipitation with Mg(OH) 2 and
CaCO 3 . In practice, treatment of hard water may not need additional
MgCO 3 , but the MgCO 3 addition would be needed for the treatment of soft
water. Individual plant water composition will need to be considered.
Coagulants [ Fe 2 (S0i 4 ) 3 , FeSO 4 , Al 2 (SO ) 3 ] removed 85% or more of the
uranium from solution containing 83 ijg/L uranium at pH 6 and 10 with a
25—mg/L dose. At pH 8, the removal efficiency was less than 50%. The pH
dependency in uranium removal appears to be related to the uranium species
and charges of coagulant at a given pH from 4 to 10. Since most municipal
water treatment plants were using water softener and/or coagulant, the
treatment pH above 10 with Fe 2 (S0 ) 3 would be more effective than pH 6
for uranium removal. Total alkalinity is the important variable and
actual removal will be water specific. The adsorbents titanium oxide
and activated charcoal also removed more than 90% of uranium in solution
at pH <7 and <7.5, respectively. Excess carbonates in the solution re-
duced the adsorption efficiency of the absorbents. Application of the
adsorbents is a possible option, but it will require additional treat-
ment costs if added to existing treatment systems.
Strong base anion exchange resin was the most effective and univer-
sal adsorbent for uranyl carbonates, which are common chemical species in
-------
11—45
natural water. Excess carbonates and polyanion species (V, Mo, etc.)
in water would reduce the removal efficiency of the resin. Although
about 10% of adsorbed uranium was remaining with resin after elution
with 1.1 N NaC1 + 0.01 N NaHCO 3 solution, the resin could be recycled
f or water treatment or washed with dilute acid after carbonate removal.
The anion exchange technology for a large volume of water treatment is
available but viii require large capital investment. Small anion
exchange columns should be considered, however, to remove uranium from
private wells located in uraniferous strata. A column containing 5 lb
of anion exchange resin could treat, for example, 106 L of water con-
taining 83 ug/L uranium with less than 1% leakage (0.8 iig/L uranium in
effluent). An efficient removal would be expected on lower uranium
waters.
The reverse osmosis system could remove uranium from the pond
water, but the removal efficiency decreased to 79% as the filtration
progressed. The results also indicated that our P0 system could treat
only about 5 L of the pond water at close to 90% removal rate. A
commercial laboratory RO unit (Cole—Partner Instrument Co., Cat. No.
C—150l—40, $1582/unit) was claimed to produce better than our results
in terms of the salt removal rate. Direct comparison of the RO perfor-
mance with other treatment methods is difficult because each of the
methods has advantages and disadvantages. However, RO will work to
lower uranium concentrations, and it is a viable option, recognizing
the large water losses involved in the reject stream. Where other
contaminants are involved, RO may be the favored option b cause of
simplicity of use.
-------
11—46
4. RECOMMENDATIONS
The results presented in this report represent bench—scale testing
on a natural water of a quality representative of uranium—bearing waters
(Ca, Mg, carbonates). The application of coagulants and/or softeners
for uranium removal is possible, generally requiring lowering or raising
the natural pH from the 7 to 8 region where uranium is most stable. The
amount of Ca, Mg, and carbonates in the feed water will influence the
magnitude of removal. For example, Table 2 of Part I showed that by
simply raising the pH of several of the municipal waters received for
uranium analysis, variable amounts of uranium were removed. This was
most likely caused by the amount of magnesium present (cannot be sub-
stantiated).
Anion exchange resins will be useful for removing uranium. On a
small scale, disposable cartridges may be employed. For reusable resin
(larger scale treatment) some problems with uranium elution are antici-
pated and will need either acid stripping or resin replacement.
-------
11—47
5. REFERENCES
Brits, R.J.N., and M.C.B. Smit. 1977. Determination of Uranium in
Natural Water by Preconcentration Anion—Exchange Resin and Delayed—
Neutron Counting. Anal. Chem. 49(l):67—69.
Grinstead, R. R., D. A. Ellis, and R. S. Olson. 1956. Recovery of
Uranium from Sulfuric Acid and Carbonate Leach Liquors by Anion Ex-
change. In: Proceedings of the International Conference on the Peace-
ful Uses of Atomic Energy, Vol. 8, United Nations, New York. pp. 45-53.
Haggag, A. M., and T. R. Stokely. 1981. Oak Ridge National Laboratory,
Oak Ridge, Tennessee. Personal communication to S. K. Hall, Oak Ridge
National Laboratory.
Hodge, V. F. 1975. Semi—Quantitative Determination of Uranium, Pluto-
nium, and Americium in Sea Water. Anal. Chem. 47(ll):1866—1868.
Kuleff, I., and K. N. Kostadinov. 1978. Epitherinal Neutron—Activation
Determination of Uranium in Environmental Waters by Np—239 After Pre-
concentration on Activated Cai’bon. J. Radioanal. Chem. 46(2):365—371.
Langmuir, D. 1978. Uranium Solution — Mineral Equilibria at Low Temper-
atures with Applications to Sedimentary Ore Deposits. Geochim.
Cosmochim. Acta 42:547—569.
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Sastri, V. S., and A. W. Ashbrook. 1976. Reverse Osmosis Performance of
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Sorg, T. J., M. Csanady, and C. S. Logsdon. 1978. Treatment Technology
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APPENDIX
Methods of Sample Preparation and Uranium Analysis
for Municpial Water Samples
Equipment
An anion exchange column was constructed by connecting a 125—mL
cylindrical glass separatory funnel with Teflon valve and tubing to a
polyethylene insert (8 mm ID, 18.5 mm long), a small capsule specially
constructed to fit inside an irradiation vial. Five holes (0.34 mm)
had been drilled into the bottom of the insert to allow passage of the
sample solution into a 250—mL filtering flask which was connected with
rubber tubing to a vacuum pump. A valve inserted into the tubing
between the filtering flask and the vacuum pump allowed the pressure to
be adjusted. Whattnan No. 50 filter paper, cut to fit the inside of the
insert, was placed on the bottom of the insert and also on top of the
resin in the insert to discourage displacement of the resin.
Reagents
An ammonium carbonate stock solution (1.3 M) was passed through an
anion exchange resin column in order to reduce any uranium contamination.
A portion of this stock solution was used to make up a wash solution
which was 0.1 /4 in ammonium carbonate and 20% in methanol. All subse-
quent dilutions were mixed from the decontaminated stock solution.
Dowex l—X8 (50—100 mesh, BioRad Laboratories, Richmond, California)
anion exchange resin in chloride form was used to preconcentrate the
uranium from the sample solutions. The resin was converted to carbonate
with 1 14 (NHL,) 2 C0 3 , stored in a solution which was 0.1 M in ammonium
carbonate and 20% in methanol. Although no uranium was detectable in
the amount of resin used, an initial acid washing step before the carbon-
ate conversion would be advisable.
Procedure
A lOO—mL portion of each sample was transferred in a disposable
beaker to a separatory funnel which had been prewashed with 0.1 M HNO 3
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and rinsed with deionized water. The sample solution was made 20% in
methanol and 0.1 M in ammonium carbonate, bringing the total solution
volume to 130 mL. The polyethylene insert was connected to the f liter—
ing flask with a small piece of tubing and filled with approximately
0.8 mL of resin. Another piece of tubing was attached to the top of the
insert and filled about three—fourths full with wash solution before
being connected to the separatory funnel. Then the valve to the separa-
tory funnel was opened and the sample solution allowed to filter through
the resin at atmospheric pressure. Air bubbles could usually be removed
by pressing on the tubing. When the flow rate slowed, a vacuum was
applied. By adjusting the stopcock in the vacuum line, a maximum flow
rate of 4 to 6 mL per minute was maintained. When all of the sample
solution had passed through the resin, the separatory funnel was rinsed
thoroughly with approximately 10 mL of wash solution. Air was then
drawn through the insert for a couple of minutes before the vacuum pump
was disconnected, and the insert was capped and dried under vacuum for a
minimum of 2 h.
To demonstrate the efficiency of the experimental setup in removing
uranium from the water, 100—mL aliqu ts of several of the water samples
received were spiked with a known amount of 7 U, the methanol and
ammonium carbonate were added, and then the sample solution was passed
through the resin column. The 237 U adsorbed on the resin was determined
by gamma counting with a thallium—activated Nal scintillation crystal
and dual photomultiplier tube detection system. A standard of the same
geometry was prepared by spiking some of the resin in an insert with the
same amount of 237 U. The water samples tested were representative of all
the samples received. The lowest recovery obtained was 73%, but a second
analysis of another lOO—mL aliquot of the same sample yielded 99%. Two
other samples showed 81 and 83% recovery, respectively, and on duplicate
runs, the recoveries were 99%. This variability in the recovery could
be due to the differences in flow rate or channeling in the resin bed,
both of which would affect the contact of the solution with the resin
beads and thus the uranium retention by the column. The inconsistencies
were therefore more likely due to handling errors rather than to being
sample—derived problems in resin sorption. To ensure meaningful results,
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most raw and final water samples were replicated, usually at an interval
of a week or so.
Controls consisting of 100 mL of deionized water instead of sample
solution were run at intervals during the analyses to ensure that the
equipment remained free of any contamination. The equipment was cleaned
with 0.1 M HNO 3 between samples and rinsed thoroughly with deionized
water. As a check on the reproducibility of the system, duplicate runs
were made on most of the raw water samples. If variability occurred
among these runs, the runs were repeated until consistent results were
obtained. Measurements of pH were also recorded for each sample.
Two problems became apparent at the beginning of the analyses.
First, Dowex 1—X2 (50—100 mesh) anion exchange resin swelled out of the
insert during the uranium preconcentration step, but a resin with greater
cross linkage, Dowex 1-X8 (50-100 mesh), was substituted and the problem
eliminated. Second, in some of the sample solutions, cloudiness developed
upon addition of the ammonium carbonate solution. A white precipitate
settled out onto the sides and bottom of the separatory funnel if the
solution was allowed to set. The solutions which precipitated were
filtered through a 0.22—pm filter before passing through the resin.
Analysis of the filter paper showed that very little uranium was incor-
porated into the precipitate.
After the preconcentration step, the insert with the uranium—loaded
resin was placed in an irradiation vial. Then, neutron irradiation of
each sample was performed in the neutron activation analysis laboratory
at the Oak Ridge Reactor located at the Oak Ridge National Laboratory.
Irradiations were performed for 60 s at a flux of 3.97 x 1013 n/cm 2 /s.
The sample was then transferred to a neutron cvunter, and counts were
taken after a delay of 6 S. Calibration was done directly with uranium
standards. The spectrum was processed with a PDP—l5 computer, using the
MONSTR program which translates the data collected into the corresponding
uranium concentrations.
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11—53
TECHNICAL REPORT DATA
i/’Icase ,rad liusrucucim on the rn cne bcforc completing)
I REPORT NO 2
3 RECIPIENT’S ACCESSIOI*NO.
EPA—57019—82—003
4 TITLE ANDSUBTITLE
S RePORT DATE
Methods of Removing Uranium from Drinking Water:
I. A Literature Survey, II. Present Municipal Water
December igg
SE MI NIZATIONCODE
Treatment and Potential Removal Methods
7 AUTHORIS)
L PERFORMING ORGANIZATION REPORT NO
J. S. Drury, D. Michelson, J. T. Ensminger, S. Y. Lee,
S. K. White, and E. A. Bondietti
9 PERFORMING ORGANIZATION NAMr AND ADDRESS
10 PROGRAM ELEMENT NO.
Information Center Complex
Information Division
11 CONTRACT/GRANT NO
Oak Ridge National Laboratory
Oak Ridge, Tennessee 37830
12 SPONSORING AGENCY NAME AND ADDRESS
Office of Drinking Water
U.S. Environmental Protection Agency
13. TYPE OF REPORT AND PERIOD COVERED
I4SPONSORINOAGENcYCODE
Washington, D.C. 20460
15. SUPPLEMENTARY NOTES
-- ancrnsrr
was searched for methods of removing uranium from drinking water. U.S.
manufacturers and users of water treatment equipment and products were also contacted
regarding methods of removing uranium from potable water. Based on the results of
these surveys, it was recommended that untreated, partially treated, and finished water
samples from municipal water treatment facilities be analyzed to determine the extent
of removal of uranium by presently used procedures, and that additional laboratory
studies be performed to determine what changes are needed to maximize the effectiveness
of treatments that are already in use in existing water treatment plants.
Uranium analyses of raw water, intermediate stage, and treated water samples from 20
municipal water treatment plants indicated that the present treatment practices were
not effective in removing uranium from raw waters when the influent concentration was
in the range of 0.1 to 16 )Ig/L uranium. Laboratory batch tests revealed that the water
softening and coagulant chemicals commonly used were able to remove more than 90% of
the dissolved uranium (<100 ug/L) in waters if an optimum pH and dosage were provided.
Adsorbents, titanium oxide and activated charcoal, were also effective in uranium
removal under specific conditions. Strong base anion exchange resin was the most effi-
cient uranium adsorbent, and an anion exchange column is a recommended option for the
treatment of private well waters containing uranium at higher than desirable levels.
17 KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
b IDENTIFIERS/OPEN ENDED TERMS
C COSATI rield/Group
Uranium
Drinking Water
08W
68D
19 DISTRIBUTION STATEMENT
Release to public
19C 1 jV 1 S (ThLS Report)
21 NO OF PAGES
20 SECURITY CLASS (ThI: page)
Unclassified
22 PRICE
EPA Form 222O 1 (9.733
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