M8 RadioChem

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3 Factors Affecting Chemical Forms of Radionuclides in Aqueous Solutions

A high proportion of samples analyzed for radionuclides consist of aqueous solutions such as natural waters and waste effluents from nuclear facilities. Solid samples need to be brought into solution before radiochemical analysis. The solutions can have a wide range of physical and chemical properties, which have an influence on the physical and chemical form of the radionuclides. Many factors affect the form in which radionuclides appear in solution. The most important of these are solution pH, redox potential, dissolved gases, ligands forming complexes with metals, humic substances, colloids, and the source and mode of generation of the radionuclides. These are discussed below. Speciation analysis, that is, analytical methods to study the forms of radionuclides, is discussed in Chapter 16.

3.1 Solution pH

The pH of the solution influences the form of radionuclides in several ways. pH directly affects the hydrolysis and redox reactions of metals, and indirectly affects the binding of metals to oxide and silicate surfaces and the formation of complexes. In this section we examine the effect of pH on the hydrolysis of metals and only briefly the other effects.

Most of the radionuclides discussed in this book are metals; notable nonmetals are carbon, selenium, iodine, and chlorine. A metal cation Mzþ hydrolyzes in water as follows:

Mzþ þxH2O $ MðOHÞz�xx þxHþ

The higher the charge on the metal and the smaller its size, the higher is the charge density of the metal and the more easily it hydrolyzes. In practice, strong hydrolysis means increasing tendency to form sparingly soluble hydroxide precipitates, and soluble hydroxide complexes and these hydrolysis products begin to form at low pH.

In aqueous solution, metal ions are surrounded by bipolar water molecules, so that the negative part of the water molecule (the oxygen, with a free electron pair) is

j35

Lehto, Jukka, and Xiaolin Hou. Chemistry and Analysis of Radionuclides : Laboratory Techniques and Methodology, John Wiley & Sons, Incorporated, 2011. ProQuest Ebook Central, http://ebookcentral.proquest.com/lib/osu/detail.action?docID=645020. Created from osu on 2019-02-26 06:54:47.

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oriented toward the positive metal ion, while the positive hydrogen atoms are directed away from it. In this way aqua complexes M(OH2)x

zþ are formed, typically with four or six water molecules in the first hydration shell. The higher the charge density of the metal, the more strongly it attracts the electron pair of the oxygen, and the bond becomes less ionic and increasingly covalent. For example, trivalent iron (Fe3þ) hydrolyzes much more easily than the divalent iron (Fe2þ) and more easily than lanthanides in the same oxidation state because the radius of iron(III) is 0.645 A

� ,

whereas that of lanthanum, for example, is much larger, 1.172 A � . Another term used

instead ofcharge density is ionicpotential, which is the ratio ofthe charge ofthe metal ion to its radius, z/r, where the radius is expressed in angstroms (A

� , 10�10 m). The

higher the ionic potential of a metal, or any element, the more strongly it attracts the electrons of the oxygen in a water molecule, and the more covalent is the chemical bond. Table 3.1 presents the ionic potentials of elements having important radionuclides.

Monovalent alkali metals, such as Csþ, do not, in principle, hydrolyze at all, nor do the heavier divalent alkaline earth metals, such as Sr2þ , Ba2þ, and Ra2þ. In contrast to this, divalent transition metal ions (Co2þ, Ni2þ, Fe2þ, Pb2þ), trivalent metal ions (Fe3þ, Ce3þ, Am3þ , Cm3þ), and metals of higher oxidation state (Po, Th, U, Np, Pu) readily hydrolyze. Raising the pH of a solution promotes hydrolysis since the aqua complexes of metals behave like weak acids, and their hydrogen ions dissociate more easily at higher pH. When a neutral hydroxide species is achieved, the metal

Table 3.1 Ionic radii and potentials of selected hexacoordinated ions of radionuclide elements. Other coordination numbers than six are shown in parenthesis (Shannon, R.D. (1976)

Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides. Acta Cryst., A32, 751).

Ion Ionic radius (A � ) Ionic potential Ion Ionic radius (A

� ) Ionic potential

Csþ 1.67 0.60 Sn2þ 1.12 1.79 Sr2þ 1.18 1.69 Sn4þ 0.69 5.80 Ca2þ 1.01 2.00 Se4þ 0.50 8.00 Ra2þ 1.48(8) 1.35 Se6þ 0.42 14.3 Ni2þ 0.69 2.90 Ac3þ 0.94 3.19 Fe2þ 0.79 2.56 Th4þ 0.94 4.26 Fe3þ 0.65 4.65 Pa5þ 0.78 6.41 Zr4þ 0.72 5.55 U4þ 0.89 4.49 Nb5þ 0.64 7.81 U6þ 0.73 8.22 Mo6þ 0.59 10.2 Np4þ 0.87 4.60 Tc7þ 0.56 12.5 Np5þ 0.75 6.67 Sm3þ 0.96 3.13 Pu3þ 1.00 3.00 Pm3þ 0.97 3.10 Pu4þ 0.86 4.65 Iþ5 0.95 5.26 Pu5þ 0.74 6.76 Cþ4 0.17 25.0 Pu6þ 0.71 8.45 Pb2þ 1.20 1.68 Am3þ 0.85 3.53 Po4þ 0.94 4.26 Cm3þ 0.85 3.53

36j 3 Factors Affecting Chemical Forms of Radionuclides in Aqueous Solutions

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hydroxide will precipitate if the amount of metal is high enough to exceed the solubility product. Figure 3.1 shows the distribution of the hydrolysis products of plutonium as a function of pH. As can be seen, the Pu4þ ion is present only in strongly acidic solution, and hydrolysis products already appear alongside it in 1 M acid, first as monohydroxide complex Pu(OH)3þ, and then, as pH increases, as di- and trihydroxide complexes, Pu(OH)2

2þ and Pu(OH)3 þ. Above pH 6, only

soluble Pu(OH)4 complexes exist in the solution. And when the concentration of plutonium is high enough to exceed the solubility product of Pu(OH)4, the hydroxide precipitates.

When the oxidation state of the metal increases to þV, not only are simple hydroxide complexes M(OH)n

5�n formed but so also are negatively charged oxoa- nions, described by the �ate� ending, and oxocations. The reason for these oxo forms is that the ionic potential of pentavalent metal ions is very high, and the positive charge of the metal attracts oxygen atoms of water molecules so strongly that all hydrogen atoms are liberated and covalent bonds are formed between the metal and the oxygen. This process takes place with nonmetals also. Elements of smaller size appear as oxoanions – among the elements having important radionuclides iodine as iodate (IO3

�) and niobium as niobate (NbO3 �). Ion sizes of the actinides, in turn, are so

large that, in pentavalent form, they bind the oxygen atoms of only two water molecules and form �yl� ions: uranyl UO2

þ, neptunyl NpO2 þ, and plutonyl PuO2

þ. Elements at oxidation state six also form �ate� and �yl� species: smaller-size

molybdenum and selenium form the negative species molybdate (MoO4 2�) and

Figure 3.1 Distribution of hydrolysis products of plutonium as a function of pH. (Choppin, G.R. (2003) Actinide speciation in the environment. Radiochimica Acta, 91, 645).

3.1 Solution pHj37

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selenate (SeO4 2�) and the larger-size actinides the �yl� cations UO2

2þ and PuO2 2þ.

Heptavalent technetium, for its part, forms pertechnetate ion TcO4 �.

Typically, radionuclides are present in solution in very low concentration and they rarely precipitate alone as hydroxides. For example, 4 MBq activity of 60Co in one liter of water is equivalent to a concentration of 1.6�10�9 M. The solubility product of Co(OH)2 is 1.6�10�15, which means that Co(OH)2 precipitates only at pH 11. Such pH values never occur in nature, where the pH rarely goes above nine. Furthermore, such high 60Co activity concentration is huge compared with the concentration found in nature. Although 60Co does not precipitate as hydroxide, it may coprecipitate with some component present in solution in macro amounts (iron, for example).

The pH of solutions also affects the oxidation states of metals, so that a decrease in pH (i.e., increase in acidity) increases the fraction of reduced forms and, conversely, a rise in pH increases the fraction of oxidized forms.

If the ligand complexing with a metal is a weak acid, such as carbonate and citrate, the pH will affect the dissociation of the acid. The ligand can form a complex with a metal only when the pH is high enough for dissociation to occur. The situation is similar in sorption of metal cations on the surfaces of hydroxides and hydrous oxides. The sorption takes place either by ion exchange through exchanging hydrogen ions of –M–OH groups on the surface or by forming surface complexes with the oxides. The –M–OH groups are weak acids, which typically dissociate only in the neutral range or in weakly acidic solution. Thus, attaching to oxides and hydrous oxides by ion exchange is possible only if the –M–OH groups have dissociated. pKa values, that is, the pH values where dissociation occurs, increase for the oxides and hydroxides typically found in natural waters in the order SiO2 < MnO2 < Fe(OH)3 < Al(OH)3.

3.2 Redox Potential

A second highly important factor for the form of an aqueous metal is the redox potential of the system. High redox potential favors oxidized forms of metals and low redoxpotential reduced forms. Manyradionuclides appear in more than one oxidation state, and because their chemical behavior depends on the oxidation state, the redox potential is of great significance. First of all, the oxidation state affects metal solubility in aqueous solutions (Table 3.2). In general, the oxidation state þIV is the least soluble and the solubility systematically increases as the oxidation state decreases from this value. Higher oxidation states are also more soluble than þIV but in a somewhat more complex manner. At oxidation state þVII all elements are in the highly soluble anionic MO4

� form, while the solubility of the þVand þVI oxidation states depends on whether the elements are in oxoanionic or oxocationic form and what are the valences of these forms. For example, the actinyl forms of actinides in oxidation state þVI (AnO22þ) are less soluble than those in oxidation state þV (AnO2

þ) because of the higher charge density of the former. Secondly, the oxidation state affects complex formation, and the stability of a metal complex follows the same trend as its solubility. As can be seen in the table below, changes in hydrogen ion

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concentration often occur in redox reactions. In systems in which oxygen participates in redox reactions, either as oxygen gas (O2) or as part of a species (e.g., IO3

�), the pH decreases in reduction reactions and, conversely, increases in oxidation reactions.

Because electrons are never free in solution, the reactions shown in the table do not occur alone. In oxidation and reduction reactions, electrons are always transferred between two components, metal or ion, as in the following reaction:

ZnþCu2þ !Zn2þ þCu

Here, zinc metal donates two electrons to copper ion, simultaneously being oxidized to oxidation state two, while the copper ion is reduced to copper metal. The reaction proceeds in precisely this direction and not the reverse because of the standard electrode potentials, E�, of the two pairs, Zn/Zn2þ and Cu/Cu2þ, participating in the reaction. For each component participating in the reaction (zinc and copper in this case) a half-reaction is defined so that the oxidized form is on the left side of the equation

Zn2þ þ2e�!Zn

Cu2þ þ2e�!Cu and for both half-reactions standard electrode potential (E�) is determined. The standard electrode potential is the potential difference, or voltage, obtained when the potential of either redox pair (e.g., zinc metal electrode in zinc sulfate solution) is measured relative to the potential of the standard hydrogen electrode. The standard hydrogen electrode consists of a platinum electrode in 1 M HCl solution which is bubbled with hydrogen gas, where the pressure of the hydrogen gas in the HCl solution is 1 atm (1.013 bar). The value of the potential of the standard hydrogen electrode when it is connected to a platinum electrode in 1 M HCl solution is set to zero. The redox half-reaction of this system is thus

2Hþ þ2e�!H2

Table 3.2 Effect of reductions on solubility of important redox-sensitive radionuclides. Reduced forms are to the right of the arrow.

Radionuclide Reduction reaction Solubility

55Fe Fe3þ þ e� ! Fe2þ Increases 99Tc TcO4

� þ 3e� þ 8Hþ ! Tc4þ þ 4H20 Decreases I isotopes IO3

� þ 4e� þ 6Hþ ! I� þ 3H2O Decreases U isotopes UO2

2þ þ 2e� þ 4Hþ ! U4þ þ 2H2O Decreases 237Np NpO2

þ þ e� þ 4Hþ ! Np4þ þ 2H2O Decreases Pu isotopes PuO2

2þ þ e� ! PuO2þ Increases PuO2

þ þ e� þ 4Hþ ! Pu4þ þ 2H2O Decreases Pu4þ þ e� ! Pu3þ Increases

3.2 Redox Potentialj39

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When the platinum electrode in 1 M HCl is replaced with another metal- or ion- containing electrode system whosepotential is to bemeasured, apotential value other than zero is obtained, positive or negative depending on whether the system is more reducing or oxidizing than the hydrogen ion system. If, for example, the potential of a zinc electrode in 1 M zinc sulfate solution is measured relative to the standard hydrogen electrode, the voltage obtained is �0.76 V. In the corresponding copper system the voltage is þ0.34 V. These voltages are the standard electrode potentials for the redox pairs Zn2þ/Zn and Cu2þ/Cu. Based on these standard electrode potentials one can predict which redox pair reduces the other: the redox pair with a more negative standard potential reduces that with a more positive standard potential. For any two half-reactions, that with the more negative standard potential moves to the left, toward the oxidized form (Zn2þ þ 2e� Zn), and that with the more positive standard potential moves to the right, toward the reduced form (Cu2þ þ 2e� ! Cu).

Standard potentials are always determined in the standard state, in which the activity of each component is one. The redox potential Eh at other concentrations of the system can be calculated from these standard potentials with the help of the Nernst equation:

Eh ¼ E�þðRT=nFÞ� lnð½ox�=½red�Þ in which R is the gas constant (8.314 J mol�1 K�1); T is the absolute temperature (K); F is the Faraday constant (96 485 C mol�1), or one electron mole; n is the number of electrons involved in the redox reaction; [ox] is the activity/concentration of the oxidized form, and [red] is the activity/concentration of the reduced form at equilibrium. If, for example, the value of the redox potential measured at 25�C for a solution containing 0.011 M iron is 0.712 V, then, with use of the Nernst equation, the ratio of the concentrations of Fe3þ and Fe2þ can be calculated as ([Fe3þ]/ [Fe2þ])¼e(((0.712V�0.771V)�nF)/RT)¼e((�0.059 � 1 � 96485)/(298 � 8.314))¼e�2.298¼0.10, where 0.771 V is the standard potential of the reaction Fe3þ þ e� ! Fe2þ. The ratio is thus 0.1; that is, 9.1% of the iron is in oxidized form and the rest in reduced form. In a redox reaction in which there are two redox pairs, for example, 5Fe2þ þ MnO4

� þ 8Hþ K 5 Fe3þ þ Mn2þ þ 2H2O, the Nernst equation is used to calculate Eh values for the two half-reactions (MnO4

� þ 8Hþ þ 5e� ! Mn2þ þ 4H2O and Fe

3þ þ e� ! Fe2þ), and their difference gives the Eh value for the whole reaction.

It also holds for redox reactions that the Gibbs free energy is DG�¼nF�E�¼ RT� lnK. The Gibbs free energy can, therefore, be calculated from either the standard potential or the equilibrium constant (K). The value of DG� indicates whether the reaction is spontaneous (negative value) or forced (positive value). As is clear from the above equation, the equilibrium constant can be calculated from the standard potential, and vice versa. Thus, if we know the standard potential, we can use it to calculate the species distribution from the equilibrium constant.

Although the standard hydrogen electrode is the normal reference electrode for measuring standard potentials, it is too inconvenient to use in measurements of

40j 3 Factors Affecting Chemical Forms of Radionuclides in Aqueous Solutions

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redox potential. The calomel electrode (Hg/Hg2Cl2) or silver electrode (Ag/AgCl) is used instead. The redox potential of a reaction is then obtained by correcting for the potential given by the calomel or silver electrode relative to that of the standard hydrogen electrode. These correction values, at 25�C, are 0.2444 V for the saturated calomel electrode and 0.19888 V for the saturated silver electrode. These values need to be added to obtain the values for a standard hydrogen electrode.

In natural systems, the ultimate determinants of redox potential are oxygen and organic matter. Soluble oxygen from the atmosphere raises the redox potential of a system, which is lowered through the aerobic decomposition of organic matter in water, soil, and bedrock. In deeper parts of the bedrock, where oxygen is depleted, and in sediment and other layers inaccessible to atmospheric oxygen, redox reactions are governed by secondary pairs, such as Fe2þ/Fe3þ, Mn2þ/Mn4þ, and S2�/SO4

2�. By acting as reducing agents for metals, microbes, too, play an important role in the redox potential. The redox potential correlates with water depth, allowing the following rough division to be made. In open surface water, the redox potential is þ200 to þ800 mV, in shallow groundwater down to about 100 meters and in deep lake and sea waters it ranges between þ200 and�200 mV, and in deep groundwater it varies between �200 and �400 mV.

Redox potential is often presented as a function of pH in Eh–pH diagrams, or Pourbaix diagrams. In the Eh–pH diagram of plutonium (Figure 3.2), the predom- inant species of plutonium are displayed as a function of Eh and pH. At the typical pH and Eh values of groundwater, Pu is predominantly present in oxidation state þIVas the soluble Pu(OH)4 complex. This exists in equilibrium with solid Pu(OH)4 if the concentration is sufficiently high. With decrease in pH, the trivalent form Pu3þ

becomes predominant. Under more oxidizing conditions, the pentavalent form PuO2

þ becomes more prevailing, and this form is also predominant in sea water. When using an Eh–pH diagram it needs to be remembered that it is primarily based on the equilibrium calculations made with initial values, such as standard potential, obtained in pure reference systems. These calculations give a good starting point for the values of each species, but, particularly for complex natural systems, practical experiments will be required to identify the real forms and their distribution. Furthermore, Eh-pH diagrams present only the predominant species in an area, and in boundary areas there are always species from both sides of the line. Species that are nowhere predominant do not appear in diagrams at all, even though they might be present in significant proportions.

Among the radionuclides, the redox potential is of greatest significance for actinides uranium, neptunium, and plutonium, which have several possible oxida- tion states. Deep in the bedrock, hundreds of meters below the surface, and in other sites with no oxygen, the redox potential is so low that it has a reducing effect not only on the actinides but on other radionuclides as well, such as those ofFe, Se,Nb, Mo, Tc, Sn, I and Po.

At the end this chapter is an appendix where are presented typical agents used to adjust oxidation states of radionuclides in radiochemical analyses.

3.2 Redox Potentialj41

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3.3 Dissolved Gases

Gases dissolved from the atmosphere play a major role in the speciation of metals in natural waters. By far the most important of these gases are oxygen and carbon dioxide.

3.3.1 Oxygen

Dissolved oxygen from the atmosphere is the ultimate determinant of the redox potential of water, and it directly affects the oxidation state of radionuclides. Divalent iron, for example, oxidizes under the influence of oxygen as follows:

Fe2þ þ1=4 O2 þ5=2 H2O!FeðOHÞ3 þ2Hþ

Surface water is continually being replenished with oxygen from the atmosphere, and the concentration of dissolved oxygen has a maximum value of 8.25 mg L�1

(0.258 mmol L�1) at 25�C and 1 bar air pressure. Oxygenated rainwater carries

Figure 3.2 Eh-pH diagram for plutonium. Plutonium concentration 0.01 mM. The upper dashed line indicates 1 atm O2 pressure and the lower dashed line 1 atm H2 pressure. These lines represent stability boundaries of water.

(Vitorge, P., Capdevila, H., Maillard, S., Faur�e, M.-H., and Vercouter, T. (2002) Thermodynamic Stabilities of MO2þx(s) (M¼U, Np, Pu, and Am), Pourbaix diagrams. J. Nucl. Sc. Technol., (Supplement 3), 713).

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oxygen deep into the soil and bedrock, but the amount of oxygen gradually diminishes with the depth. Oxygen is consumed first and most intensively in the oxidation of organic matter in the top layer of the soil.

RCHOþO2 !CO2 þH2O

(RCHO represents the organic matter, e.g., cellulose) Still more oxygen is consumed lower down in the soil profile by organic matter

transported there by water. Oxygen in soil and bedrock is also consumed by microorganisms, which use it in energy production. The oxidation of organic matter and microbial activity are intimately related. Other ways in which oxygen is con- sumed are in the oxidation of metals, for example, the oxidation of iron(II) to iron(III) and manganese(II) to manganese(IV), and the oxidation of sulfides to sulfates. In unsaturated soil (soil that is not saturated with groundwater), oxygen diffuses more deeply into the ground relatively fast, but maximally at the rate it diffuses in air (0.205 cm2/s). The oxygen consumed in the oxidation of organic matter is then replenished by fresh oxygen from the atmosphere. The diffusion rate of oxygen in water is over ten thousand times slower than that in air. Thus, the oxygen consumed inoxidation inwaterlogged soilis replenished much more slowlythan thatconsumed in unsaturated soil, and the redox potential is lower in waterlogged soil. And if oxygenated water does not seep through fractures in the bedrock, the redox potential will decrease with increasing depth. Oxygen is also efficiently consumed by the organic matter in water: just 3 mg L�1 of dissolved organic matter can consume the maximal amount of oxygen that can dissolve in water. Hundreds of meters deep in the bedrock, where spent fuel from nuclear power plants will be buried, there is practically speaking no oxygen at all and the redox potential in ground water will be extremely low, as low as �400 mV.

The oxygen in rivers, lakes, and the sea is used up in the same way as that in the soil, that is, mostly in the oxidation of organic matter. If there is no vertical flow of water, as may happen in frozen lakes in winter, oxygen may become depleted and anoxic conditions develop in the bottom water. In anoxic conditions, Fe(OH)3 and MnO2 precipitated under oxidizing conditions may dissolve, with Fe3þ reducing to soluble Fe2þ and Mn4þ to soluble Mn2þ. Simultaneously, any radionuclides coprecipitated with them will be remobilized.

3.3.2 Carbon Dioxide

The other major gas important for the speciation of metals in water is carbon dioxide. Carbon dioxide reacts in water to form carbonic acid: CO2(aq) þ H2O K H2CO3. The first dissociation constant (pK1) of carbonic acid, in the reaction H2CO3 K H

þ þ HCO3

�, is 6.35, and the second (pK2), in the reaction HCO3 � K Hþ þ CO32�, is

10.33. Thus, hydrogen carbonate ion is predominant when the pH exceeds 6.33 and carbonate ion when the pH exceeds 10.33 (Figure 3.3). Carbonic acid is the predominant acid in natural waters and the primary determinant of pH of natural

3.3 Dissolved Gasesj43

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waters, while the carbonate ion, CO3 2�, effectively precipitates many metals and

forms complexes with them. Carbon dioxide is dissolved into water from the atmosphere, where it makes up

0.033 volume percent. The concentration is also expressed as partial pressure, which in a gas mixture, in this case the air, is equal to the mole fraction multiplied by the total pressure. At one bar pressure, at 25�C, the partial pressure of carbon dioxide in dry air is thus 10�3.5 bar (0.00033 bar). The partial pressure of carbon dioxide in natural waters generally ranges between 10�4 and 10�2 bar. In surface water, carbon dioxide is almost always in equilibrium with atmospheric carbon dioxide, and its partial pressure is the same as that in the atmosphere. Aquatic plants take up carbon dioxide for photosynthesis, and, at times of vigorous growth, the partial pressure of carbon dioxide will drop below the equilibrium level if the atmosphere is unable to restore the balance quickly enough. Typically, carbon dioxide builds up in the bottom water of lakes and the seas as a consequence of the decomposition of aquatic plants and the oxidation of organic matter in the bottom sediment. Carbon dioxide is also introduced into the bottom water through the influx of carbonate-rich groundwater. Increase in the amount of carbonic acid intensifies the anoxic bottom conditions.

In soil, the amount of carbon dioxide is determined by the aerobic decomposition of organic matter in the unsaturated zone (RCHO þ O2 ! CO2 þ H2O), which continually adds carbon dioxide to the soil. During warm growing periods, the partial pressure of carbon dioxide in the soil may rise as high as 0.1 bar. Part of the generated carbon dioxide dissolves in water, and part escapes from the soil to the atmosphere. If the ground freezes in winter and the carbon dioxide is unable to escape to the

Figure 3.3 Carbonate species in water as a function of pH.

44j 3 Factors Affecting Chemical Forms of Radionuclides in Aqueous Solutions

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atmosphere, the amount of carbon dioxide in the soil may increase dramatically. Carbon dioxide, and dissolved organic material in general, are carried with water to deeper levels in the soil. The amount of carbon dioxide is greater in deeper, anoxic soil and sediment layers and in the bedrock than in surface layers because of the many processes generating carbon dioxide (reduction of sulfate and nitrate, formation of methane) and the inability of carbon dioxide to escape from deep in the ground to the atmosphere. Even though the amount of carbon dioxide is high in groundwater, the pH of the water increases rather than decreases because of reactions that consume hydrogen ions (e.g., the reduction of sulfate, 2CH2O þ Hþ þ SO42�K 2CO2 þ HS� þ 2H2O, where CH2O represents organic material). As a result, the pH of deep groundwater is distinctly higher than that of shallow groundwater, ranging between 8 and 9.

In surface water, as well as soil water and groundwater, carbonate can precipitate as calcium carbonate, CaCO3, usually in the form of calcite, if the concentrations of calcium and carbonate are high enough for the solubility product to be exceeded. The solubility product of calcite, Ks¼ [Ca2þ ]� [CO32�], is 3�10�9. Along with sodium, potassium, and magnesium, calcium is one of the four cations most commonly dissolved in natural waters in macro amounts. In lakes the concentration of calcium is typically 0.25–1.25 mmol L�1 and in the sea 10 mmol L�1. Calcite will also precipitate in soil and bedrock and in bottom water of lakes when the amount of carbon dioxide is sufficiently high.

The relevance of all this to the speciation of radionuclides is that many metal radionuclides can coprecipitate with calcite. Most of them form sparingly soluble carbonates (see Table 4.1). Alkali metals are exceptional in not forming sparingly soluble carbonates and do not coprecipitate with calcite. Strontium is closely related to calcium in chemical behavior, and 90Sr, for example, strongly coprecipitates with calcite. One of the methods for speciation analysis that is introduced later (Chapter 16) is selective/sequential extraction, which is used to determine the speciation of radionuclides in soil and sediment samples. A key step in these extractions is to determine the radionuclides bound to carbonates. The carbonates are dissolved in acetic acid at a pH of about 4, which releases radionuclides associated with them. Calcite is least soluble at a pH greater than 9 but dissolves even in weak acids such as acetic acid.

Besides precipitating with calcite, carbonate affects the chemistry of radionuclides in natural waters by forming strong complexes, especially with the actinide elements, whose formation constants with carbonate are higher than those of most other ligands present in water (sulfate, phosphate, chloride). Typically fluoride forms even stronger complexes, but its concentration in natural waters is very low compared to that of carbonate. The main competitor to the formation of carbonate complexes is the formation of hydroxide complexes by hydrolysis. Particularly noteworthy in the formation of carbonate complexes is the behavior of uranium in groundwater, where uranium appears in oxidized form as uranyl ion UO2

2þ. In water where the pH exceeds 6, uranyl ion forms highly soluble negatively charged carbonate complexes: UO2(CO3)2

2� at pH 6–7 and UO2(CO3)3 4� at pH above 7. The other actinides also

form strong carbonate complexes. The strongest complexes are formed by tetravalent

3.3 Dissolved Gasesj45

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actinides. The stability of carbonate complexes of the actinides typically decreases in the following order of actinide ions:

An4þ > AnO2þ2 > An 3þ � AnOþ2

Figure 3.4 shows the speciation of neptunium as a function of pH at a total CO2 concentration of 2.3 mmol L�1, respectively. At low pH (<7.5), Np is mainly as NpO2

þ, while at higher pH, NpO2OH, NpO2CO3 �, and NpO2(CO3)2

3�are the major species of neptunium.

3.4 Ligands Forming Complexes with Metals

The ligands in water available for complex formation markedly affect the chemical forms of radionuclides, since most of the metal cations present in natural waters in trace concentrations appear in complexed form. Complexes are formed with inor- ganic and organic anions and with neutral organic molecules. The most important of the complexes formed with organic molecules in water are humic complexes, which are discussed in the following section. Of the inorganic ligands, the most important are OH�, CO3

2�, HCO3 �, SO4

2�, HPO4 2�, PO4

3�, Cl�, and F�. Of these, hydroxide complexes were covered earlier in the context of hydrolysis and carbonate complexes in the discussion of dissolved carbon dioxide. Table 3.3 shows the concentrations of inorganic ions in fresh water and in sea water. The concentrations in fresh water are clearly lower than those in sea water, and they vary over a wide range, while the concentrations in sea water are more or less constant.

As described earlier, in pure water solution, metal cations are present in the form of aqua complexes, typically surrounded by four or six water molecules in the first hydration shell closest to the metal. The above-mentioned inorganic ligands form both outer-sphere and inner-sphere complexes with metal ions. Outer-sphere com- plexes form between hydrated metal cations and negative ligands with no change taking place in the first hydration shell. The bond is purely electrostatic, which means that it is an ion–ion interaction. Because there are water molecules between the metal and the ligand, the metal and the charged ligand lies far apart from one another. In inner-sphere complexes, in turn, the ligand replaces one or more of the water molecules in the first hydration shell and approaches the metal, causing the covalent character of the bond to increase. The stability of the complex increases at the same time. The move from ionic to covalent bond occurs gradually, and the bonds of many complexes have both ionic and covalent character. The nature of the bond is determined by the metal and ligand charges and their size, that is, their charge densities.

In natural waters, outer-sphere complexes are mostly formed with the ions Naþ, Mg2þ, and Ca2þ , which occur in macro concentrations. The fourth important macro cation, potassium, does not form complexes at all. The outer-sphere complexes of sodium, magnesium, and calcium are generally called ion pairs. The weakest complexes are formed with hydrogen carbonate, chloride, and sulfate, and elements

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Figure 3.4 Aqueous speciation diagram of neptunium at 25�C. Total Np concentration 0.01 mmol/L, total carbonate concentration 2.3 mmol/L.

3.4 Ligan

ds F orm

in g C om

plexes w ith

M etalsj

Lehto, Jukka, and X iaolin H

ou. C hem

istry and A nalysis of R

adionuclides : Laboratory T echniques and M

ethodology, John W iley &

S ons,

Incorporated, 2011. P roQ

uest E book C

entral, http://ebookcentral.proquest.com /lib/osu/detail.action?docID

= 645020.

C reated from

osu on 2019-02-26 06:54:47.

other than alkali and alkaline earth metals also form outer-sphere complexes with these ligands. Among the stronger complexing ligands are fluoride, hydrogen phosphate, and especially carbonate, with which transition elements and higher charged metals form inner-sphere complexes.

Actinides form their strongest complexes with carbonates, and when the pH is above 6 these are the predominant species; only hydrolysis products compete with them. The strength of actinide complexes decreases in the order carbonate > hydrox- ide > fluoride�hydrogen phosphate > sulfate > chloride. Complex forms are often presented in graph form as a function of pH at defined total concentration. As an example, uranium phosphate and hydroxide complexation as a function of pH is presented in Figure 3.5.

3.5 Humic Substances

Humic substances are decomposition products of plants formed in the organic layer of the soil, and they are also carried downward, with water, into groundwater. Humic substances form when the primary polymeric decomposition products of plants – cellulose, hemicellulose, and lignin – decompose further in oxic conditions. Humic substances are classified into fulvic acids, humic acids, and humin on the basis of their molecular size and solubility. Fulvic acids are smallest in size, with a nominal molecular weight of 500–1500 g mol�1. They are also most soluble, being soluble in both acidic and alkaline solutions. Humins are the largest of the humic substances, with molecular weights up to about 5000 g mol�1. Humins are insoluble in both acid and alkaline conditions. Humic acids fall between fulvic acids and humins: their molecular weight is several thousand grams per mol, and they are soluble in alkaline but not in acidic conditions (pH < 2). Since the pH of natural waters never falls below pH 2, humic acids are always present there in soluble form. Soil contains other decomposition products of plants as well: organic components of small molecular size, such as carbohydrates and proteins. These differ from humic substances chemically, in their precisely determinable compositions and structures, and thus their chemical behavior can be described with thermodynamic formulations. Humic substances, in contrast, are polymerized organic materials of variable composition

Table 3.3 Anion concentrations in fresh water and sea water (Stumm, W. and Morgan, J. (1995) Aquatic Chemistry, 3rd edn, John Wiley & Sons, p. 290).

Ion Fresh water Sea water

HCO3 � 0.1–5 mmol/L 2.5 mmol/L

CO3 2� 0.001–0.1 0.03

SO4 � 0.01–1 28

Cl� 0.01–1 550 F� 0.001–0.1 0.06

48j 3 Factors Affecting Chemical Forms of Radionuclides in Aqueous Solutions

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Figure 3.5 Uranium phosphate and hydroxide complexes as a function of pH. Uranium concentration 0.084 mM, phosphate concentration 2.6 mM.

3.5 H u m ic S u bstan

cesj 49

Lehto, Jukka, and X iaolin H

ou. C hem

istry and A nalysis of R

adionuclides : Laboratory T echniques and M

ethodology, John W iley &

S ons,

Incorporated, 2011. P roQ

uest E book C

entral, http://ebookcentral.proquest.com /lib/osu/detail.action?docID

= 645020.

C reated from

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and molecular size containing acidic groups, namely carboxyl groups (R�COOH) and less acidic phenol groups (�OH group joined to an aromatic ring). Because phenolic acid is very weak, the phenol groups in humic substances are not dissociated in natural waters with a pH in the range 4–9 and phenol groups are far less important than carboxylic groups in binding metals. Humic and fulvic acids sorb metal ions not only through their carboxylic groups but also through the free electron pairs of oxygen and nitrogen atoms in their structure. Because of the high surface area of humic substances, the ion exchange capacity of the organic layer of soil is high: 100–200 meq/100 g. Figure 3.6 presents a hypothetical picture of humic acid.

Formation constants can be determined for the metal complexes formed between most radionuclides and inorganic and organic ligands, such as carbonate and citrate. The formation of complexes between radionuclides and humic and fulvic acids is less straightforward because their structures and the numbers and types of functional groups participating in complex formation vary. The behavior of radionuclides in natural waters, however, can be modeled with use of apparent formation constants determined for the complex formation with humic and fulvic acids of defined composition.

Humic and fulvic acids are washed from the soil into surface water, and are also formed there in the decomposition of aquatic plants. The concentration of humic and fulvic acids together is about 0.1 mg L�1 in groundwater, 0.1–1 mg L�1 in seawater, and 1–10 mg L�1 in rivers and lakes. Concentrations are greatest in mire lands, bogs, and marshes, and in water draining from these: tens or even a hundred milligrams per liter. Humic and fulvic acids strongly bind radionuclides, especially those with a high charge, and thus their affinity for the actinides is especially high. Actinides in complexes of humic and fulvic acids in natural waters nevertheless appear mostly in trivalent form (Am3þ, Pu3þ) because hydroxide ions compete more strongly than humic and fulvic acids for tetravalent actinides (Th4þ, U4þ, Np4þ, Pu4þ), and carbonate complexes are predominant with penta- and hexavalent actinides (UO2

2þ, NpO2

þ, PuO2 þ, PuO2

2þ). Fulvic acids, which are the most soluble part of humic material, play a key role in

leaching metals downwards in the soil. By dissolving and forming complexes with iron and aluminum and with trace metals, including radionuclides, fulvic acids carry

Figure 3.6 Model structure of humic acid (Stevenson, F.E. (1982) Humus Chemistry, Wiley- Interscience, New York).

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these down into the mineral layers of the soil. With time, the fulvic acids break down through oxidation, and the bound iron and aluminum (and radionuclides), are released and precipitate as oxides into the enrichment zone of the soil.

3.6 Colloidal Particles

Another important factor besides humic and fulvic acids that influences the speciation of radionuclides in natural waters is the association of radionuclides to inorganic colloidal particles. Colloidal particles are the fine particles in water varying in diameter from a nanometer to several hundred nanometers. Though 0.45 mm is often taken as the upper limit, the definition is not precise. Colloidal systems are relatively stable, or metastable, whereas larger, suspended particles will sediment out within a reasonable time. Colloidal particles are stable, or persistent in solution, because their rate of diffusion due to thermal motion is higher than their rate of sedimentation due to gravity. As the size of particles increases, the rate of sedimen- tation begins to exceed the rate of diffusion. The lower size limit for colloidal particles is similarly not precise. Thermodynamically, colloidal particles are distinguished from dissolved components in that chemical potential can be determined for dissolved components but not for colloidal particles. An exact composition can be assigned to truly dissolved matter, and the behavior of dissolved components can be described with conventional thermodynamic concepts. Although in many systems colloidal particles will appear to behave like dissolved matter, their true behavior differs in an essential way. If colloidal and dissolved forms of the same element are present in a system, the behavior of the two forms will be distinct. Their separation from each other is nevertheless often very difficult.

Colloidal particles abound in natural waters; there may be millions of them in one milliliter, and their total surface area is high. The composition of colloids varies widely. The humic and fulvic acids discussed in the previous section are important organic colloidal particles. Most inorganic colloidal particles are clay minerals, that is, aluminum silicates and hydrous oxides of iron, manganese, titanium, aluminum, and silicon. OH groups on the surface of the oxides are largely dissociated in the neutral pH range of natural waters, allowing metal cations to adsorb by ion exchange. Metal cations, especially those with highcharge,can also attachto thesurface through complexation, that is, formation of covalent bonds with surface oxygen atoms. Radionuclides not only attach to clay minerals through surface OH groups but also absorb in the interlayer space in the crystal lattice, which results in a strong bonding. Cesium sorption by clay minerals is an important example of the latter process. Though cesium is in general a very soluble element, sorption by clay mineral particles, if present, results in effective sedimentation of cesium into lake bottoms by the clay minerals and makes cesium rather immobile in clayey soils.

Colloidal particles containing radionuclides are classified into two groups. The first group is the carrier colloids, which include humic and fulvic acid colloids and colloids formed with oxides onto whose surfaces radionuclides have sorbed. Also included in

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this group are particles in which radionuclides are coprecipitated, isomorphically or nonisomorphically, along with the compound of some closely related element. The second group is the intrinsic colloids, or colloidal particles containing a single radionuclide, typically in hydroxide or oxide form. The formation of intrinsic colloids is rare because the concentrations of radionuclides in natural waters are generally very low, and the solubility product is seldom exceeded. Only a few radionuclides are ever present in solution in such large amounts that the solubility product is exceeded. In nature, only uranium and thorium appear in such large amounts, and even then the formation of intrinsic colloids is highly unlikely, especially in case of uranium, which is very soluble in natural waters.

3.7 Source and Generation of Radionuclides

The forms of radionuclides are also affected by their source. The major sources of radionuclides in the environment are nature�s own processes, nuclear weapons tests in the atmosphere, accidents at nuclear power plants, nuclear weapons production facilities, fuel reprocessing plants, nuclear power plants and nuclear waste repositories. These have already been discussed in Chapter 1. Here we briefly focus on the effect of the formation process on the physical and chemical form of radionuclides.

The most important source of radionuclides is nature itself, most radionuclides being members of the decay series of uranium and thorium. Among these nuclides are U, Th, 226Ra, 222Rn, 210Pb, and 210Po. With the notable exception of radon, they are almost entirely bound in minerals present insoil andbedrock, and to some degree appear in dissolved form in ground waters and surface waters. Radon occurs in gaseous form and also diffuses into the atmosphere, from which its daughters (210Pb, 210Po, and short-lived lead, polonium and bismuth isotopes) deposit on the earth�s surface attached to aerosol particles. Within the soil there are also a few individual long-lived natural radionuclides, the most notable being 40K. A large group of lighter, cosmogenic radionuclides are generated in the atmosphere. Of these, 14C and 3H are mostly present in the gas phase, while many others, such as 7;10Be, are attached to aerosol particles and precipitate onto the Earth�s surface either as dry deposition or, in rain or snow,as wet deposition. Also, industrial plants spread natural nuclides in the environment. For example, the burning of coal and peat releases natural nuclides into the atmosphere, and smelters and fertilizer plants discharge natural nuclides in soluble form to watercourses and the sea.

A second important source of radionuclides is the nuclear weapons tests carried out in the atmosphere in the 1950s to 1970s. Of all the fallout radionuclides associated with these tests, the major ones still present in the environment are 90Sr, 137Cs, 238;239;240;241Pu, and 241Am. Most of the fallout was global stratospheric fallout, in which the radionuclides that deposited on the ground were bound to atmospheric aerosols. Over the years, the radionuclides have depositedonto the surfacelayer ofthe soil. At actual test sites (Nevada, Semipalatinsk, etc.) and the areas round about,

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however, a major part of the radionuclides, especially of plutonium and americium, is present as large, insoluble particles, some consisting of unfissioned plutonium and uranium. Most of the transuranium elements at test sites are present on the soil surface.

A third source is accidents at nuclear power plants, the most consequential of these being the explosion and fire at the Chernobyl plant in 1986. Most of the nuclides were the same as those released to the environment in nuclear weapons tests, but there were also notable differences. Unlike the debris from nuclear weapons tests, the fallout from Chernobyl rose in an emission plume and was carried over Europe in the troposphere at a height of 1000–2000 meters. It was not only aerosol particles with adsorbed radionuclides that were swept over Europe but also larger, hot particles – fuel fragments and condensation products of weakly vaporizing metals – where radionuclide concentrations were high. These hot particles are distinctly less soluble than the particles in global fallout from nuclear weapons tests. The proportion of hot particles in the fallout from Chernobyl was greatest closest to the accident site and decreased with distance. The nuclide composition and isotope ratios also differed from those of weapons fallout. Thus, the fallout from Chernobyl contained, in addition to the fission product 137Cs and also the activation product 134Cs, which was not found in deposition from weapons tests. Likewise, the ratios of plutonium isotopes in the fallout from Chernobyl and that from weapons tests were noticeably different, allowing these ratios to be used in source identification.

The emissions from nuclear power plants are very low compared with the previous two sources. Small amounts of 137Cs, 60Co, 54Mn, and 63Ni are discharged to rivers and the sea, and settle in sediments in the near surroundings. Tritium remains in the water phase. 14C and 85Kr may be released to the atmosphere, where they are diluted in air masses.

Over the decades, nuclear weapons production facilities and civilian fuel reprocessing plants have released high amounts of radioactive materials to the environment, and in particular to surface water. Particularly polluted are areas around nuclear weapons production complexes in Russia, the worst being the area around the Majak nuclear complex in the southern Urals. In the past two or three decades, effluents from the reprocessing facilities at La Hague in France and Sellafield in Great Britain have been reduced to a fraction of what they once were, but still a moderate amount of radioactive material is discharged to the sea.

Large amounts of radionuclides end up to the environment in the final disposal of nuclear waste. At present, only low and intermediate active waste generated during the operation of nuclear plants is disposed of in final repositories, which means that the waste mostly contains relatively short-lived fission and activation products, such as 137Cs, 60Co, and 63Ni. In future, waste from decommissioned nuclear plants will be taken to repositories. This includes waste such as radioactive concrete, where the nuclide composition is much the same as that in the waste generated in plant operations. The most active nuclear waste that will be brought to repositories will, however, be the spent nuclear fuel and the waste generated in reprocessing, both containing a broad spectrum of short- and long-lived fission and activation products, such as 90Sr, 99Tc, 135Cs, 137Cs, and isotopes of plutonium and 241Am. From these

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final repositories, the nuclides will find their way into the biosphere tens of thousands of years from now, at the earliest, and then only small amounts of those that are longest lived.

3.8 Appendix: Reagents Used to Adjust Oxidation States of Radionuclides

Radionuclide analyses often include adjusting the oxidation state of the radionuclide. A number of reagents can be used for this purpose, of which some of the most important are briefly described here.

3.8.1 Oxidants

Hydrogen peroxide H2O2 is an oxidizing agent in acidic medium, the oxygen in it having an oxidation state of �I and aiming at a more stable oxidation state of �II. Hydrogen peroxide, for example, oxidizes tetravalent technetium to pertechnetate in the following way:

2TcO2 þ3H2O2 !2TcO�4 þ2H2Oþ2Hþ

H2O2 is also often used as an oxidizing agent to decompose organic matter to convert all radionuclide species into their inorganic forms.

Oxyanions such as NO3 �, Cr2O7

2�, S2O8 2�, ClO�, MnO4

�, and ClO4 � are

oxidizing agents. However, their oxidizing powers vary greatly. The last two, permanganate and perchlorate, are the strongest oxidizers. Perchloric and nitric acids are not often used to adjust the redox states of radionuclides; rather, they are widely used to decompose organic matter by oxidation to drive the radionuclides bound in the organics into solution for analysis. The use of perchloric acid, although very efficient, is mostly avoided since it can result in an explosion if the fumes are not flushed properly. Permanganate ions are highly oxidizing and can be used to oxidize Np(V) to Np(VI), for example. Permanganate is so oxidizing that it can result even in the reduction of hydrogen peroxide to oxygen while the permanganate is reduced to manganese dioxide:

2MnO�4 þ3H2O2 !2MnO2 þ3O2 þ2H2Oþ2OH�

Nitrite NO2 � is used as NaNO2 to adjust the oxidation state of plutonium to þIV.

It is very useful since it can both oxidize trivalent plutonium to the tetravalent state and, at the same time, reduce hexavalent plutonium to the same tetravalent state when Fe3þ is present as a catalyst. Since plutonium can exist in several oxidation states at the same time, nitrite is an important agent when only tetravalent plutonium is needed. The reduction of hexavalent plutonium to tetravalent is a slow process since it requires Pu�O bond breakage. The rate of this reaction, however, can be enhanced by addition of Fe3þ.

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Pu3þ þNO�2 þ2Hþ !Pu4þ þNOþH2O

PuO2þ2 þNO�2 þHþ !Pu4þ þNO�3 þOH�

3.8.2 Reductants

Metal ions at their less stable lower oxidation states, such as Ti3þ, Cr2þ, Fe2þ , can be used to reduce radionuclides from their higher oxidation states to lower ones. An example is the reduction of tetravalent plutonium to trivalent with Ti3þ.

Ti3þ þPu4þ !Ti4þ þPu3þ

Ascorbicacid (vitamin C) C6H8O6 is usedto reducetrivalent ironto its divalent state. Trivalent iron interferes with many radionuclide analyses, as in Am3þ determina- tions, for example.

2Fe3þ þC6H8O6 !2Fe2þ þC6H4O6 þ2Hþ

Sulfamic acid H3NSO3 or ferrous sulfamate Fe(NH2SO3)2 is often used to reduce plutonium to Pu3þ . It has been used, for example, in the separation of plutonium from uranium in the PUREX process in the reprocessing of nuclear fuel; however, while it has now been replaced by other reductants, it is still used in analytical radiochemical separations.

Sulfurous acid H2SO3 and sulfite Na2SO3 or Na2S2O5 (becoming NaHSO3 when dissolved in water) are reductants used in radiochemical separation processes. As they can rapidly reduce the high valence states of Pu (Pu4þ, PO2

þ, PuO2 2þ) to Pu3þ

and Np (NpO2 þ, NpO2

2þ) to Np4þ, they are often used for this purpose. For example, in the separation of Pu from large volumes of water, Pu is first co-precipitated with Fe(OH)3. Since the recovery of Pu in the form of PuO2

þ and PuO2

2þ is very low, Pu is reduced with sulfite to Pu3þ, for which the recovery is good. In chromatographic separations of Pu, all oxidation states of Pu should be adjusted to Pu(IV) because of its high affinity to chromatographic columns (strongly basic anion exchange resin or TEVA Resin). This is normally carried out in two steps: first, high oxidation states of Pu (PuO2

2þ ,PuO2 þ, Pu4þ) are reduced to Pu3þ by sulfurous acid

(or sulfite in acidic solution), and then Pu3þ is oxized to Pu4þ by nitrite (or by using NO2/NO2

� in concentrated nitric acid). In addition, sulfurous acid is also often used as a reductant in the separation of 129I from water, since it can rapidly reduce all high oxidation states of iodine (IO4

�, IO3 �, IO�, I2) to I

�; I� is then oxidized to I2 with nitrite for extraction to the organic phase. Since sulfite is a strong but unstable reductant, especially in an acidic condition, the water solution of sulfite is normally freshly prepared before utilization.

Hydroxylamine hydrochloride NH2OH�HCl is an often used reductant to reduce elements to their lower oxidation state. In radiochemical separations of Pu, it is often eluted from an anion exchange column or an extraction chromatographic column (for example TEVA) by reducing Pu(VI) with NH2OH�HCl to Pu3þ in dilute

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hydrochloric acid solution. Hydroxylamine is, however, not a strong reductant: it can, for example, reduce IO3

� to I2 but not to the lowest oxidation state of iodine (I �).

Hydrazine N2H4 or a hydrazium salt (for example hydrazine sulfate, N2H4�H2SO4) is also a reductant and can be used to reduce Pu to Pu3þ. The reducing ability of hydrazine is similar to that of hydroxylamine. The reduction rate of Pu to Pu3þ by hydrazine is slower than that using sulfamic acid; however, the stability of hydrazine and hydroxylamine is much better than that of sulfamic acid and sulfite, especially in acidic solution.

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