**7. Uranium complexation with an emphasis on phosphorus**

Uranium complexation pairs a central cation (coordination center) with a surrounding array of molecules and ions. Phosphorus interactions with U(VI) have been studied to assess whether phosphorus may reduce the availability and mobility of U(VI) [12, 65−67]. Stojanovic et al. [18] reported that phosphorus may readily form uranyl phosphates and subsequently precipitate autunite. They noted that at pH levels greater than 6.0, the dominant U(VI)-phosphorus species was the plant-available UO2 PO4 species, whereas at more acidic soil reactions, UO2 HPO4 and UO2 H2 PO4 + were more abundant and are not considered as plant-available U-phosphate species. Grabias et al. [65] studied uranyl acetate immobilization in ferruginous soils amended with phosphates. In acidic pH ranges, a strong U(VI) sorption was observed in the presence of phosphate, supporting their premise that adsorption was promoted by the formation of UO2 (H2 PO4 )(H3 PO4 )+ , UO2 (H2 PO4 )2 and (UO2 ) 3 (PO4 )3 4H2 O.

Mehta et al. [67] demonstrated that U(VI) flux in soil columns was substantially reduced when phosphate was present. Sequential extractions demonstrated that the U(VI) could be readily extracted by ion-exchange and dilute acid treatments. Laser-induced florescence spectroscopy inferred adsorption to be the dominant retention mechanism.

Sandino and Bruno [68] determined the solubility of (UO<sup>2</sup> )3 (PO4 )2 4H2 O (s) and the formation of U(VI) phosphate complexes over the pH range of pH 4–9. In their study, UO2 HPO4 and UO2 PO4 − were the dominant U species. Minimum U(VI) solubility for the (UO2 )3 (PO4 ) 2 4H2 O (s) system occurred near pH 6, whereas the minimum U(VI) solubility for amorphous (noncrystalline) and crystalline schoepite occurred near the pH levels of pH 7.4 and 8.4, respectively. Thermodynamic data for U(VI) with respect to phosphate and carbonate from the literature are well-documented by Sandino and Bruno [68].

Lenhart et al. [69] described uranium(VI) complexation with citric acid, humic acid and fulvic acid in acidic media (pH 4.0 and 5.0). Using Schubert's ion-exchange method, the U(VI)-citric acid complex was determined to be 1:1 uranyl-citrate complex (β1,1 = 6.69 ± 0.3 at I = 0.10). Humic and fulvic acids were demonstrated to strongly bind to U(VI), with humic acid forming a slightly stronger binding complex. The U(VI)-humic acid and U(VI)-fulvic acid complexes were determined to be non-integral (1 U(VI) with between 1 and 2 humic or fulvic acids), suggesting that a 1:1 stoichiometry involving a limited number of high-affinity sites.

Ivanov et al. [70] observed uranyl sorption on bentonite in the presence of humic acid with trace levels of uranium(VI). Uranyl sorption on bentonite was shown to be strongly pH dependent. In the absence of humic acid, U(VI) sorption exhibited a sorption edge between pH 3.2 and pH 4.2. In the presence of humic acid, U(VI) sorption slightly increased at low pH and curtails at moderate pH. Soluble uranyl carbonate species inhibited U(VI) sorption at alkaline pH levels. At pH intervals from pH 3 to pH 4, UO2 HA was predicted ([U] = 8.4 × 10−11 and pCO2 = 10–3.5 atm, HA = humic acid). From pH 5 to pH 7, UO<sup>2</sup> (OH)HA was predicted to be the dominant species. Tinnacher et al. [71] studied the reaction kinetics of tritium-labeled fulvic acid on uranium(VI) sorption onto silica, demonstrating that metal sorption rates are a complex function of metal and organic ligand concentrations and the nature and abundance of mineral surface sites.

Sandino and Bruno [68] reported the oxalate and sulfate complexation reactions involv-

PO4

+ H+ 1.69 ± 0.15

exudates supported greater uranium concentrations in the adjacent soil solution. Sandino and Bruno [68] provided phosphate complexation reactions involving the uranyl cation: (1) UO<sup>2</sup>

The MinteqA2 simulation of U(VI) at 10−3 mol U/L demonstrated that the dominant U(VI)-

**pH 4 pH 6 pH 8**

Sulfate, log β = 1.92. Tandy et al. [72] reported that citrate and malate from root

at I = 0 (Grenthe et al. [55]).

2+ + PO4

3− = UO<sup>2</sup>

**PO4**

(30.6%). Rutherfordine and (UO<sup>2</sup>

at pH 4 and 6, whereas at pH 8, the dominant species

Oxalate, log β = 6.02 and (2) UO<sup>2</sup>

Chemical Thermodynamics of Uranium in the Soil Environment

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PO4

2+ +

131

2+

1−, log β = 13.25 ± 0.09.

) 3 (PO4 ) 2 were

2+ + Oxalate2− = UO<sup>2</sup>

(OH)2 + 2H+ log K = −5.4

CO3 log K = 14.11

)2 log K = 48.61

)2 log K = 46.00

)PO4 log K = 25.00

2+ 0.76 ± 0.15

+ H+ 1.12 ± 0.07

)2 + H+ 0.87 ± 0.05

Ka1, Ka2 and Ka3 constants are (−2.14 ± 0.03), (−7.21 ± 0.02) and (−12.35 ± 0.03), respectively.

, log β = 7.28 ± 0.10 and (2) UO<sup>2</sup>

Additional equilibrium constants are presented in **Tables 7**–**9**.

**8. Simulation of uranium complexation with H3**

4− (67.9%) and UO<sup>2</sup>

(HPO4 )2

> (HPO4 ) 2

UO2 5.3 (3.1%) 7.3 14.4

(OH) 7.2 7.2 12.3

(OH)2 8.2 8.2 18.4

(OH)<sup>5</sup> 11.5 7.5 18.9

ing the uranyl cation: (1) UO2

HPO4

Sulfate2− = UO<sup>2</sup>

2− = UO<sup>2</sup>

phosphate species were UO2

formed as solid phases (**Table 10**).

**Species −log (activity)**

(CO3 ) 3

+ HPO4

UO2 2+ + 2H2

UO2 2+ + CO3

2UO2

2UO2

UO2 2+ + H+ + PO4

UO2 2+ + H3

UO2 2+ + H3

UO2 2+ + 2H3

UO2 2+ + 2H3

H3 PO4 O = UO<sup>2</sup>

2− = UO<sup>2</sup>

PO4 = UO<sup>2</sup>

PO4 = UO<sup>2</sup>

PO4 = UO<sup>2</sup>

PO4 = UO<sup>2</sup>

3− = Ca(UO<sup>2</sup>

3− = Fe(UO<sup>2</sup> ) 2 (PO4

3− = H(UO<sup>2</sup>

H3 PO4

H2 PO4 +

> (H3 PO4 )H2 PO4 +

(H2 PO4

**Table 9.** Experimental equilibrium data for the U(VI)-H3

)2 (PO4

**Table 8.** Precipitation reactions involving U(VI) (Chen and Yiacoumi [40]).

**Reaction log β**

2+ + Ca2+ + 2PO4

2+ + Fe2+ + 2PO4

were UO2

UO2

(UO2 )2

(UO2 )3


**Table 7.** Formation constants for selected aqueous species (Davis [44]).


**Table 8.** Precipitation reactions involving U(VI) (Chen and Yiacoumi [40]).

Mehta et al. [67] demonstrated that U(VI) flux in soil columns was substantially reduced when phosphate was present. Sequential extractions demonstrated that the U(VI) could be readily extracted by ion-exchange and dilute acid treatments. Laser-induced florescence spec-

> )3 (PO4 )2 4H2

O (s) and the formation

)3 (PO4 ) 2 4H2 O

HA was predicted ([U] = 8.4 × 10−11

(OH)HA was predicted to

HPO4

and

troscopy inferred adsorption to be the dominant retention mechanism.

of U(VI) phosphate complexes over the pH range of pH 4–9. In their study, UO2

were the dominant U species. Minimum U(VI) solubility for the (UO2

(s) system occurred near pH 6, whereas the minimum U(VI) solubility for amorphous (noncrystalline) and crystalline schoepite occurred near the pH levels of pH 7.4 and 8.4, respectively. Thermodynamic data for U(VI) with respect to phosphate and carbonate from the

Lenhart et al. [69] described uranium(VI) complexation with citric acid, humic acid and fulvic acid in acidic media (pH 4.0 and 5.0). Using Schubert's ion-exchange method, the U(VI)-citric acid complex was determined to be 1:1 uranyl-citrate complex (β1,1 = 6.69 ± 0.3 at I = 0.10). Humic and fulvic acids were demonstrated to strongly bind to U(VI), with humic acid forming a slightly stronger binding complex. The U(VI)-humic acid and U(VI)-fulvic acid complexes were determined to be non-integral (1 U(VI) with between 1 and 2 humic or fulvic acids), suggesting that a 1:1 stoichiometry involving a limited number of high-affinity sites.

Ivanov et al. [70] observed uranyl sorption on bentonite in the presence of humic acid with trace levels of uranium(VI). Uranyl sorption on bentonite was shown to be strongly pH dependent. In the absence of humic acid, U(VI) sorption exhibited a sorption edge between pH 3.2 and pH 4.2. In the presence of humic acid, U(VI) sorption slightly increased at low pH and curtails at moderate pH. Soluble uranyl carbonate species inhibited U(VI) sorption at

be the dominant species. Tinnacher et al. [71] studied the reaction kinetics of tritium-labeled fulvic acid on uranium(VI) sorption onto silica, demonstrating that metal sorption rates are a complex function of metal and organic ligand concentrations and the nature and abundance

CO3 + 2H+ log K = −7.01

<sup>+</sup> log K = 0.30

<sup>−</sup> log K = −6.35

CO3 log K = −16.68

2− + 4H+ log K = −16.43

4− + 6H+ log K = −28.45

<sup>−</sup> + 5H<sup>+</sup> log K = −17.54

Sandino and Bruno [68] determined the solubility of (UO<sup>2</sup>

130 Uranium - Safety, Resources, Separation and Thermodynamic Calculation

literature are well-documented by Sandino and Bruno [68].

alkaline pH levels. At pH intervals from pH 3 to pH 4, UO2

of mineral surface sites.

O + H2

CO3 = UO<sup>2</sup>

CO3 = UO<sup>2</sup>

CO3 = UO<sup>2</sup>

<sup>−</sup> = UO<sup>2</sup> NO3

2− = HCO<sup>3</sup>

2− = H<sup>2</sup>

CO3 = (UO<sup>2</sup>

(CO3 )2

(CO3 )3 )2 CO3 (OH)3

**Table 7.** Formation constants for selected aqueous species (Davis [44]).

2UO2

UO2 2+ + H2

UO2 2+ + 2H2

UO2 2+ + 3H2

UO2 2+ + NO3

H+ + CO3

2H+ + CO3

2+ + 3H2

and pCO2 = 10–3.5 atm, HA = humic acid). From pH 5 to pH 7, UO<sup>2</sup>

UO2 PO4 −


**Table 9.** Experimental equilibrium data for the U(VI)-H3 PO4 at I = 0 (Grenthe et al. [55]).

Sandino and Bruno [68] reported the oxalate and sulfate complexation reactions involving the uranyl cation: (1) UO2 2+ + Oxalate2− = UO<sup>2</sup> Oxalate, log β = 6.02 and (2) UO<sup>2</sup> 2+ + Sulfate2− = UO<sup>2</sup> Sulfate, log β = 1.92. Tandy et al. [72] reported that citrate and malate from root exudates supported greater uranium concentrations in the adjacent soil solution. Sandino and Bruno [68] provided phosphate complexation reactions involving the uranyl cation: (1) UO<sup>2</sup> 2+ + HPO4 2− = UO<sup>2</sup> HPO4 , log β = 7.28 ± 0.10 and (2) UO<sup>2</sup> 2+ + PO4 3− = UO<sup>2</sup> PO4 1−, log β = 13.25 ± 0.09. Additional equilibrium constants are presented in **Tables 7**–**9**.

#### **8. Simulation of uranium complexation with H3 PO4**

The MinteqA2 simulation of U(VI) at 10−3 mol U/L demonstrated that the dominant U(VI) phosphate species were UO2 (HPO4 )2 at pH 4 and 6, whereas at pH 8, the dominant species were UO2 (CO3 )3 4− (67.9%) and UO<sup>2</sup> (HPO4 ) 2 (30.6%). Rutherfordine and (UO<sup>2</sup> ) 3 (PO4 ) 2 were formed as solid phases (**Table 10**).



Typically, the pH range of minimal U(VI) mineral solubility coincides with the pH range for optimal U(VI) adsorption. U(IV) complexes are frequently less soluble and less mobile than U(VI) complexes [73]. Duquene et al. [23] noted that U(VI) reduction to less soluble U(IV), by either biotic or abiotic processes, influenced uranium mobility. Stojanovic et al. [17] confirmed that soil temperature, pH, oxidation–reduction potentials and the presence of complexing agents were important factors influencing uranium bioavailability and plant uptake. Shahandeh and Hossner [33] employed a selective sequential extraction protocol to show that U(VI) partitioned into exchangeable, carbonate, manganese, iron, organic and residual fractions. In soils where the carbonate fraction was expected to be important, appreciable plant uptake of U(VI) into the roots and culms of a wide variety of plants was demonstrated. In soils having U(VI) partitioning into iron, manganese and organic fractions, the U(VI) plant uptake was substantially smaller.

In a review, Langmuir [38] reported solution U(VI) speciation data from pH 7 groundwater at Yucca Mountain (Nevada, USA) with a total U(VI) concentration of 10−8 mol/L. The U(VI)

at 7.9%, (2) UO<sup>2</sup>

used a double layer model for simulating U(VI) adsorption on a smectite (montmorillonite).

+

Uranium(VI) may be adsorbed onto Fe-oxyhydroxides which may subsequently pursue dis-

especially in alkaline solutions at elevated pH levels. Surface properties of soil mineral phases have altered chemical's reactivity because of the presence of small quantities of noncrystalline Fe- and Al-oxyhydroxides. Thus, these alterations of chemical affinity may be attributed to differences in surface area, abundance and composition of Al-OH, Fe-OH and Si-OH groups, and

<sup>+</sup> 12.3 −0.95

+ H+ 7.1 0.15

(OH)<sup>5</sup> + H+ −15.8 −16.80

at 0.06% and (6) UO<sup>2</sup>

(CO3 ) 2

PO4

of 2.60, (6) AlO-(UO2

)3 (PO4 )2 4H2

Chemical Thermodynamics of Uranium in the Soil Environment

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at 83.1%, (3) UO<sup>2</sup>

)3 (OH)<sup>5</sup>

**Log K for ≡Al Log K for ≡Si**

of −9.73, (2) > Al(OH)<sup>2</sup>

O(s) = 3UO<sup>2</sup>

(CO3 ) 3

of 8.33, (3) > SiO−

of −14.95 and (7)

S or (2) desorbed,

at 0.8%. Pabalan and Turner [57]

+

2+ +

133

at 7.8%,

Sandino and Bruno [68] provided the solubility estimate for (UO<sup>2</sup>

CO3

tinctive pathways: (1) U(VI) undergoes reduction to U(IV) by mobile Fe2+ or H2

other features that impact the structure of adsorption surfaces (**Table 11** and **12**).

SOH = SO− + H+ −13.6 −6.95

**Table 11.** Adsorption site reactions and surface protonation/deprotonation reactions (McKinley et al. [46]).

(OH)2

of 2.70, (5) > SiO-UO<sup>2</sup>

Their surface complexation constants were (1) > AlO<sup>−</sup>

+

of −15.29.

O as log Kso ± 2σ = 48.48 ± 0.16.

2PO4

(4) UO2

SiO-(UO<sup>2</sup>

SOH + H<sup>+</sup> = SOH<sup>2</sup>

2+ = SO-UO<sup>2</sup>

) 3 (OH)<sup>5</sup> +

S is the surface site representing Al and Si.

<sup>+</sup> = SO-(UO<sup>2</sup>

) 3

SOH + UO<sup>2</sup>

SOH + (UO<sup>2</sup>

3− + 4H2

**10. Uranium adsorption**

percentage speciation was: (1) UO2

F at 0.007%, (5) UO<sup>2</sup>

of −7.20, (4) AlO-UO<sup>2</sup>

)3 (OH)<sup>5</sup>

Total U concentration was 10−3 mole/L, which was allowed to equilibrate and allow precipitation of rutherfordine and (UO2 )3 (PO4 )2 .

Activity coefficients were determined by the Debye-Huckel equation.

Calcium and H3 PO4 concentrations were initially standardized at 10−3 mole/L. The presence of CO2 (g) at 2 × 10−2 bar (2 kPa) and an ionic strength standardized by 0.01 *M* NaNO3 . Within a pH column, ( ) indicates the percentage of the U species.

**Table 10.** The MinteqA2 simulation of U(VI) species in the presence of CO2 (g) and H3 PO4 .
