Solubility of Actinides in Various Solutions: A Review of Studies on Th, U, Np, Pu, and Am, Study notes of Literature

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APPENDIX G

SUMMARY OF LITERATURE SOLUBILITY STUDIES

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-2.

-4.

-5.

X

• m *mil .k:x i

+0.6M KCI. 7 days, Felmy et el. (1991) X0.6M NaCI, 7 days, Felrny at al. (1991) x0.8M NaCI. 98 days, Felmy et at. (1991) -1.2M NaCI. 7 days, Felmy et at. (1991)

• 1.2M NaCI, 98 days, Felmy et el. ( l m l ) 03M NaCI, 8 days, Felrny et al. (1991) A3M NaCI, 122 days, Felmy et al. (1991) m3M NaCI, 372 days, Felmy et at. (1991)

.I

m+&

I a*

• A ?'

' A

m A .

rn

X

X

4

x x

I X

1 7 8 9 10 11

-10.

2 3 4^5^6

Figure Q-

Hydrous Th(lV) Oxlde Solublllty as a Functlon of pH for dllute and Brine Solutlons

1

• 2 3 4 5 6 7 8 9

10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

6 27 28 29 30 31 32 33

35 36 37 38 39 40 41 42 43 44 45 46 47 48

U (VI). The presence of organic materials and corrodable metals in the repository environment will probably restrict U to the IV state.

Figure G-2 summarizes U solubility data for uraninite (UO,). Solubility data reported by Parks and Pohl (1988) for UO, at 100 and 150°C in dilute NaCl solution (0.1 M) indicate a decrease in the concentration of U from -7 to -9.5 log M as pH increases from 1 to 7. Gray (1986) investigated the solubility of unirradiated UO, fuel-rod pellets in a saturated NaCl brine of pH 6.2 to 6.4 and reported U concentrations of -7 to -8 log M, about 1.5 to 2.5 orders of magnitude lower than UO, solubility values in dilute NaCl solutions. The solubility data of Gray (1986) serve as a conservative estimate of U concentrations in Salado brine, because the solubility of UO, increases with temperature and the solubility data were obtained at elevated temperatures, relative to the ambient temperature of about 30°C in the WlPP repository horizon.

Figure 6-3 summarizes U solubility data for schoepite (U03*2H,0) in HCIO, solutions at 25°C.

Krupka et al. (1985) obtained solubility data that indicate a decrease in the U concentration from -2 to -5 log M as pH increased from 3 to 9. This trend was followed by an increase of similar

magnitude in the U concentration over the pH interval of 9 t o 12. Bruno and Sandino (1989)

conducted solubility experiments with amorphous and crystalline UO,-2H,O, which show the amorphous schoepite to be one to two orders of magnitude more soluble than the crystalline form. Solubility studies for schcepite in saturated NaCl brine were not found. However, based on the

studies presented on Figure G-2 for UO,, it is expected that schoepite solubility in Salado brine

will be 1 to 2 orders of magnitude greater than the dilute solution data on Figure G-3.

4.0 NEPTUNIUM

The dominant valence states for Np in the natural environment are IV, V, and VI. As is the case with U and Pu, Np solubilities are highly dependent on valence state. Np (IV) is expected to dominate in the reducing repository environment, with solubility being controlled by crystalline NpO,. Other forms of Np (IV) such as amprphous NpO,, NpO,*xH,O, and Np(OH), will age to crystalline NpO,. An important Np(V) solubility-controlling phase is Np0,OH.

Figure 6-4 summarizes Np (IV) solubility data for NpO, and NpO,*xH,O in dilute solutions and

NaCl brines. The studies of Rai and Strickert (1980) and Rai et al. (1982) at 25'C show a

decrease in Np concentration from -5.5 to -6.5 log M as pH increases from about 2 to 6. The

data of Kim et al. (1985b) were obtained in 1M and 5M NaCl solutions at 20 to 25 ' C. The runs

at pH of 5 to 6 resulted in Np concentrations of about 4.5 log M, and those at a pH of 7.5 to 8 show a Np concentration of about -6 log M. Ionic strength does not appear to affect the solubility of Np in these 1 M and 5M runs, but a comparison of these data over the pH range of

5 to 6 with the data of Rai and Strickert (1980) suggests the solubility of NpO, is increased by

about 1.5 orders of magnitude in NaCl solutions of moderate to high ionic strength.

Pryke and Rees (1987) investigated the solubility of NpO,*xH,O in solutions equilibrated with concrete at ambient temperature. Results on Figure G-4 show an apparent decrease in the Np concentration of about one-half log unit as the pH increase from 10 to 13. Comparing the studies of Pryke and Rees (1987) and Rai and Strickert (1980), it appears that the data of Pryke and Rees (1987) anchor the projected trend of the data of Rai and Strickert (1980).

A U 0 8 9 5 M I P f l A C B S R G G-3 753435.01 10/12/95 5:33prn

.. .'

-2. 0

(^0 )

8

' 0 0 0 4 0 (^4 4) 0

0 0 B 0 0 4^0

0 0

0cp0%

%

0 0 0

b 0 O

a

OU0,.2Hz0 Krupka el el. (1985) 4UO32H2O (a) Bruno 8 Sandlno (1989) OU0$2Hz0 (c) Bmno 8 Sandlno (1989)

0 8

"'B

0 I 2 3 4 5 6

E

1 -4. a -

9 v,

-5.

-8.

Figure (3- U03:2H20 Solubllity as a Function of pH for Dilute Solutions

-4.

-5.

-6.

s 8

8

=

• -7.

9 0 ) il

-8.

i f-,, /

c

I c ' j

-9.

Figure G-

Np02 Solubility as a Function of pH for Dilute, Saline, and Brine Solutions

ONpO,, Rai et el. (1982)

ONp02, Rai and Strickert (1980)

ANp0,:xH20. Pryke and Rees (1987)

m237Np02, 1M NaCI, Kim et ai. (1985b)

A237Np02,5M NaCI. Kim et el. (1985b)

m A A

0 0 0

(^0) O O o

A A

0

B 0

A

• A A AA A A A A A A

- - - - - - - - - - x - - - A & A A. _ O_* A t o (^) n - 03 00

4. 0 L _ O 0_*

A.. " O

A AA

A s

A.

A

*0~~~Np0,0H(a).Neck et al. (1992) uNp0,OH. €wart et el. (1985) 0 ~ ~ ~ N p 0 ~ 0 H ( a ) .1M NaC104, Neck et el. (1992) 237NpOpi, 1M NaCl04, Klm et al. (1985b) m237Np0,0H, 1 M NaC104, Neck et al. (1992) Az37Np020H,3M NeC104, Neck et al, (1992)

~ ~ ~ ~ N p 0 , 0 H ,5M NaCI. Klm et al. (1985b)

--237NpOz0H, 1M NaCI, Kim et al. (1985b)

0

e

-

6 7 8 9 10 11 12 13 14

PH

Figure 6- Np020H Solublllty as a Functlon of pH for Dilute, Saline, and Brine Solutlons

-

0239Pu(OH)3(a),Rai et al. (1987)

o3ePu(OH)3(a),Ral et al. (1987)*

0239Pu(OH),~a),PBB,, Rai etal. (1967)

A239Pu(OH)3(,), PBB,, Rai el at. (1987)

0 0 0

0 0 0 0 (^0) Q 0

PH

Figure (3-

Amorphous ~ ~ * P u ( O H ) ~Solublllty as a Functlon of pH for Dllute Brlne and Solutlons

I 3 4 5 6 7 8 9

11 12 13 14 15 16 17 18 19 20 21 22 23 24

• ! 27 28 29 30 31 32 33

48 -* 9

is about two orders of magnitude greater in the brine relative to the dilute solution used by Ewart et al. (1985).

Figure G-8 summarizes solubility studies carried out by Kim et al. (198s) and Kim et al. (1985b) with usPuO, in dilute solutions and NaCl brines. The solubility of 238Pu0,in a 0.1M NaCl solution decreases sharply from -5.5 to -7 log M over the pH interval of about 3 to 3.5, followed by an apparent increase in Pu concentration to about -6 log M as the pH rises to 5.5. Data points for the 3M and 5M NaCl brines indicate an increase in Pu concentration as pH increases. These observed trends are the opposite of most actinide solubility trends, which show decreasing actinide concentrations with increasing pH. Kim et al. (1985a) attributed the enhanced solubility to radiolysis effects and colloid formation.

Solubiiity data for p9PuOzare summarized on Figure G-9. The data reported by Rai et al. (1980)

on u9Pu0, solubility in dilute solutions shows a decrease in Pu concentration from approximately -

5 to -9 log M as pH increases from about 3 to 8. Solubility data reported by Kim et al. (1985b)

for pePuO, follow a decreasing trend that is parallel to the data of Rai et al. (1980) but at Pu

concentrations that are one to two orders of magnitude lower over the pH interval of 4 to 6.5. The large difference in the solubility of p9Pu0, obtained by these two independent studies may be due to variation in the degree of crystallinity of the z"PuO, solid used in the experiments. Additional solubility studies with uoPuOpwere carried out by Kim et al. (l985b) using 5M NaCl solutions, and these data plot below the dilute solution data obtained by Kim et al. (1985b) between pH 3.5 and 5 but then converge around a pH of 7.

Data on the solubility of PuO,*xH,O and paPuOz(OH), in dilute solutions are summarized on

Figure G-10. Pryke and R e e s (1987) examined the solubility of PuO,*xH,O in cement-equilibrated

solutions adjusted to selected pH values with HCI. Their data show a sharp decrease in Pu

concentration from about -6 to -10 log M as pH increases from 7 to 9, followed by a shallow

decrease as pH rises to 12. Solubility data for Pu (Vl) was investigated by Kim et al. (1985b) using 2"P~0,(OH),. The data of Kim et al. (1985b) indicate a decrease in the Pu concentration from about -4 to -8.5 as the pH increases from about 5.5 to 10. Above a pH of 10, the solubility trend for usP~O,(OH), shows a slight increase in Pu concentration.

6.0 AMERICIUM

Americium can be present in the natural environment in the 111, IV, and V valence states. The valence state with the largest stability field is the 111 state.

Figure G-11 summarizes data on the solubility of Am(OH),. Rai et al. (1983) investigated the solubility of amorphous 241Am(OH),in dilute solutions adjusted with HCI or NaOH to set pH, and the solubility of^ amorphous 2*3Am(OH),^ in a pure H,O^ solution that had pH adjusted with HCI or

tetrapropyl NH,OH. These data sets indicate a decrease in Am concentration from approximately

-4 to -10 log M as the pH increases from about 7 to 10. Above a pH of 10, the trend of the plotted solubility points is essentially flat, and the Am concentration is maintained at around -10 log M.

Kim et a1 (1985b) examined the solubility of Am(OH), in a 0.1M NaCIO, solution and their data follow a trend similar to the data of Rai et al. (1983). but for a given pH the solubility data of

--
AL/OB-95/WPEACBS'R3744.G G-11 763435.01 10/12/95 5:Sprn

-5.

-8.

c

! J

-7.

G) A h)

-8.

Figure Q-

23Pu02 Solubility as a Functlon of pH for Dilute and Brine Solutions*

0 0 0 0 0

. A

0 8 0 A

0

0 ..^ +

0.

. o~~PuO,,Kim*^ et^ at. (l985a)

0238Pu02,Klm et at. (1985b)

Az38Pu02.3M NaCI, Klm et al. (1985a)

+238Pu02,5M NaCI, Kim et al. (1985a)

e23BPuOz,5M NaCI, Klm et al. (1985b)

-5.0 0 0 0

m

0 ono 0

0

0 0

OPuO,:xH,O, Pryke and Rees (1987)

023BPu0,(OH)2,Kim et al. (1985b)

0 0 O 0000 0 0000

00

0

-11.0 1

(^5 8 7 89) I 6 1 1 12 13 14

pH

Flgure 0-

Pu02. x H20 and PuOp (OH), Solubilities as a Functlon of pH for Dllute Solutions

-3.

-4.

-5.

-6.

-7. E

cj^8

5 -8.

• c?.

-9.

-10.

-11.

-12.

024'Am(OH)3(a),Ral et al. (1983) 0243Am(OH),(a~,Rai et al. (1983) AZ4'Am(OH),. Kim et al. (1965b)

XAm(OH),, Ewart et al. (1985)

OAm(OH),, Ptyke and Rees (1987) mAm(OH),, sat NaCl, Flambard et al. (1966)

5 6 7 (^8) 9 10 11 12 13 14

PH

Figure 0-1 1

Am(OH)3 Solubility as a Function of pH for Dilute and Brlne Solutions

?sol07 crl.M.W.Q%d A?

I

I

-4.

0 -. -l

-8.

A

II M A

A

- A

0

o ~ ~ ~ A ~ O H C O , .Felmy et al. (1980) OAmOHCO,, Pryke and flees (1987)

024'Am020H,5M NaCi, Kim el at. (1985b)

A2*'Am02, Kim et el. (1985b)

A241Am02,5M NaCI, Kim el al. (1985b)

, 1

0 0

OO 0 (^0 ) 0 0

0 O 0 n

-9.0 I

3 4 5 6 7 8 9 10 11 12 13

PH

Figure G-

AmOHC03, Am020H, and Am02 Solubilities as a Function of pH for Dilute and Brine Solutions

7eol07 a3.W W. W A l 0

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