Application of environmental colloid science in the soil systems

Adsorption and its Applications in Industryand EnvironmentalProtection Studies in Surface Science and Catalysis,Vol. 120 A. Dabrowski(Editor) 9 1998Elsevier Science B.V. All rights reserved.

351

Application of environmental colloid science in the soil systems J. Szczypa a, I. Kobal b, and W. Janusz a aDepartment of Radiochemistry and Colloid Chemistry, UMCS, 20-031 Lublin, Poland b Josef Stefan Institute, Jamowa 39, 1000 Ljubljana, Slovenia and School of Environmental Sciences, Vipavska 13, 5000 Nova Gorica 1.

INTRODUCTION

Soil is, beside air and water, the most important environment for the living organisms including h u m a n beings. Between these systems, some components exchange, influencing the development and condition of organisms. Up till now most of the h u m a n food comes directly or indirectly from soil cultivation. That is the reason for the special care of soil system quality. The main way is to fully understand complex processes that rule the ecosystem. When air and water systems are simple, the soil is very complicated and variable one, concerning the mineral, grain-size distribution and composition of surrounding solution. Soil is multicomponent, polydispersed system, usually treated as a three-phase system (solid, liquid and gas), although some authors consider organisms the fourth phase. Tens thousands of soils can be distinguished. They can differ by their provenance, parent minerals and age. For example later soils, less weathered, are rich in silica, alumina and iron minerals whereas old weathered ones are without many soluble minerals of alumina and iron and consists of silts and rusty oxides [1]. Looking at composition of soil there can be distinguished primary and secondary minerals. Quartz, orthoclase, plagioclase, muscovite, biotite, pyroxenes and olivine are in the first group and hydrolyzed silicon oxides, aluminum and iron hydroxide, carbonates (calcite and dolomite), hydroxides and silt minerals are in the second one. Half of the most soils is formed by minerals, the rest consists of water solution, air and organic substances (about 5%). In sandy silts the organic substance contents is small and reaches few percent whereas in peat soils or muck the organic substance contents may reach 100% [1]. Among many organic substances in soil very important role play humines, humine acids, humatomelitic acids and fulvic acids. These compounds dissolve in water in a different degree. The solid particles of soil show various size from bigger t h a n l m m treated as gravel and stones, then sand 0.05-lmm, silt 0.002-0.05mm to clays < 0.002mm. The sandy soil, formed by fine particles is concise, plastic, sticky and impermeable. This

352 type of soil is defined as heavy soil. The soil with sand is porous and defined as light soil. In this soil the transportation processes of air and water with nutrient substances are easier. The presence of humus provides proper structure of soil, decides degree and mechanism of soil particle aggregation. Adsorbing on mineral particles, h u m u s provides spongy structure of soil that promotes a transportation process. Soil is a source of heavy metals and their adsorbent also. These factors, which have an influence on the total contents of metals and their amount assimilated by organisms is essential for h u m a n beings life and soil fertility. Some of heavy metals are known as microelements (Cu, Zn, Co, Mn) that are indispensable for growth and their life at specific concentrations and toxic at higher, and other metals that are potentially toxic elements (PTE) (As, Hg, Pb, T1 and U) [2]. Radioactive isotopes, existing in the environment may be divided according to their origin: primary, cosmic and anthropogenic ones. Primary radioisotopes, which came into being during nuclear synthesis of elements have a half-life time comparable or bigger than the age of our planet. To this group belong K-40 and isotopes of U-235, U-238 and Th-232 series. The isotopes of cosmic origin form in atmosphere of the Earth as a result of cosmic radiation for example H-3 and C-14. Man has inconsiderable influence on the formation and spreading of the isotopes from these groups. Unfortunately there is a group of isotopes that were introduced to environment by man. Some of them, by the occasion of nuclear energy production (from mining of radioactive ores to closing down the exploited reactors), other as a result of nuclear weapon tests. The main radioisotopes introduced to environment in this way include the uranium family nucleus' fission products (U, Pu) and products of activation of reactor materials. The list of radioactive isotopes, important for the environment contains elements of the most groups of the periodic table. The list of elements and their origin was presented by Lieser [3]. A physico-chemical behavior of the most radioactive isotopes in soil is the same as their stable equivalents, only for H-3 some differences may be distinguished. Radioisotope concentration may be similar or considerably smaller than stable isotope in environment, depending on its half-life time. The concentration of the hydrolyzable element is connected with a form it exists in the environment. When the concentration is smaller than the solubility product of the respective hydroxide or other insoluble salt and there are no conditions to adsorption on the dispersed particles of other phase, the isotope may form non-ideal solution. When the isotope concentration is so high that exceeds the solubility product then the dispersed phase may form. When the dispersed phase exists and physical conditions enable the adsorption on the solid, then isotope forms pseudocolloid, that means exists in the dispersed phase not forming crystalline structure. From the chemical and physical behavior at the solid/solution interface there is no difference between stable and a radioactive isotope. The only difference is their harmfulness to living organisms. To fully understand the transport phenomena of nutrient substances in such complicated system, the adsorption measurements of the selected element on the

353 soil are not sufficient method. It is necessary to learn surface properties of individual components of soil and on this basis to work out the model of incorporated processes. There are some papers that treat processes running in soil as a physical chemistry of dispersed systems that is ion adsorption and colloid transportation in water [2-7]. In present paper the essential processes, having influence on the ion distribution between solid and solution are presented. They may be related to soil minerals/ aqueous solution system and thus may be useful for the understanding of the transportation and adsorption phenomena in such complex system as the soil is. In the subsequent chapters of this paper the problems of the ion adsorption and surface charge formation are presented.

2.

METAL OXIDE/AQUEOUS SOLUTION SYSTEM

Among metal oxides found in soil, beside silica, there are oxides, hydroxides and hydrated oxides of iron, aluminum and mixes hydroxides. The electrolyte ions may accumulate on the metal oxide surface as a result of nonspecific adsorption, caused by electrostatic interaction, complexing, ion exchange. Another mechanism of accumulation of ions on the soil minerals is heterocoagulation of colloids which have been formed by ions. The adsorption process, connected with the influence of coulombic interaction, is defined as nonspecific adsorption. It is caused by the electric charge on the surface of the metal oxide, which caused distribution of the ions in the surrounding solution layer. In consequence, the ions of the same sign as the sign of surface charge of the particle of the solid will be removed from that region of the solution, whereas the ions of the opposite charge will be accumulated. The layer of the solution that is under the influence of coulombic forces of the solid surface charge is called the diffusion layer. The whole range of the solid with accumulated charge, together with the part of solution with compensating charge is defined as an electrical double layer (edl). The distribution of the charge in the diffusion layer is relatively well described by Gouy-Chapman theory of the edl, where the charge density near the flat plane of the solid is equal:

~d = - ~ / 8 ~ ~

sinh/F~d)2RT

(I)

where: ~d - diffuse layer charge density, ~- relative dielectric constant for water = 78.25, ~o-absolute dielectric constant = 8.85*10~2CemU-~, c - electrolyte concentration, R - g a s constant - 8.314 J*mol~*K-~, T - temperature, ~gd- diffuse layer potential, F - Faraday constant 96500 C'mole 1.

354 Electric charge on the metal oxide surface is formed because of acid-base reactions of the surface hydroxyl groups. In colloid chemistry there are two approaches to an electric charge formation on the surface of the oxide. The first one by the reactions of the ionization surface group, defined as 2-pK model [8]: -SOH~

~

-SOH+H

(2)

§

- S O H +_~ - S O - + H +

(3)

where - S represents some surface of metal oxide irrespective of metal atom n u m b e r that coordinates the oxygen atom of the hydroxyl group. The reaction 2 and 3 constants may be calculated from the charge density data as a function of pH or from ~ potential versus pH dependence. The alternate attitude to the formation of the charge on the metal oxide is proposed by MUltiSIte Complexation model (MUSIC model) where hydroxyl group on the oxide surface are gifted by the charge that depends on the degree of saturation of the oxide valence by coordinating metal atoms and hydrogen, connected by hydrogen bound by donor or acceptor bonding [9]. (--MekO(H)m(HH20)n~ init + H + ~

( - M e k O ( H ) m + l ( H H 2 0 ) n - l ~ fin

(4)

where: Sinit(fin) - charge of the surface group before or after adsorption of hydrogen, which depends on the number and charge of coordinated metal atoms (k and SMe) and number (m) of hydrogen atoms (donor type connection) and (n) number of hydrogen atoms (acceptor type connection), s = ~ SMr + m* s H + n* (1 - s~) + V, SH = 0.8, whereas V = -2. For the number of free orbitals of the oxygen there is restriction for the hydrogen bonds for k = 1 ~ n + m = 2, for k = 2 ~ n + m =2 or n + m = 1 whereas for triple coordinated oxygen atoms (k = 3), n + m = 1. That means that hydrogen atom may be bound to surface oxygen in the acceptor or donor way [9]. Reaction constants of surface groups in MUSIC model are calculated theoretically from crystallographic data. First, from Brown theory, the valence of metal in the lattice of metal oxide is calculated [10], next, the charge of the surface group and finally, on this basis, the constants of the surface groups. The increase of the electrolyte concentration in the metal oxide/electrolyte system causes the increase of the charge density at the interface due to the following reactions: - SOH~An-SOH+Ct

~

+ ~

- SOH + H § + An-

(5)

-SO-Ct ++H +

(6)

According to the site binding theory, anions, which reacts with the hydroxyl group, produce surface complex type compound. The positive charge of this group is

355 located in the surface layer. that occupies position in the concentration of -= SOH ~Ansolution and increase with Similarly, for reaction 6, the

It is compensated by the negative charge of the anion inner Helmholtz plane (IHP). Following reaction 5, the groups should decrease with the increase of pH of the the increase of the concentration of the electrolyte. adsorption of the cation results in the formation of the

surface compounds - S O - C t +- type complex, where negatively charged part is in surface plane of edl, whereas the cation is in the IHP. Because the ions, adsorbing according to the reactions 5 and 6, form the complex type connections not only by electrostatic but also chemical forces, this type of adsorption is called specific adsorption of the ions [11]. The Cs-137 and Cs-134 isotopes adsorb in the same way as the stable caesium. However, because of their low concentrations (for example 1Bq/dm 3 Cs-137 equals to 1"10 -15 mole/dm 3) towards 1:1 salts present in soil water, the processes of specific and nonspecific adsorption will have minor importance. The excess of monovalent cation Na § or K § will lower the adsorption of Cs § because of the competitive adsorption on the same site. The concentration of potassium ions, in aqueous solution of the soil, reaches 2mg/dm 3 (about 50 ~mole/dm~). These ions will adsorb on hydroxyl groups according to reaction 6, whereas the adsorption of Cs § ions from such dilute solution will be limited [12]. For example complexation constant of Na § for TiO2, pKNa=8.2 and Cs § pKc~=7.2. For the concentrations Na+=lmmole/dm ~ and Cs+=lpmole/dm 3 [lkBq of Cs-137/dm 3] the [-TiO-Na ~] to [-TiO-Cs § relation, calculated with neglecting of the radius difference and activity coefficient was 10-s, based on the following equation:

I TiO-Cs TiO-Na-

Kcs * [Cs+]= 10 -7,2 ,10-12 KNa [Na + ]

= 10 -8

(7)

10 .8,2 , 1 0 - 3

That confirms the above opinion of the negligibly low specific adsorption of the Cs-137 or Cs-134 from the soil solutions containing other also alkaline metal cations Na § or K +. Only contamination by Rb-87, whose concentration (1Bq/dm 3 =1 mole/dm~), may cause appreciable specific adsorption. After all, caesium isotopes may adsorb on metal oxides by the exchange reaction with respective ions, presented in the oxide for example as contamination. Another process, responsible for the deposition of the caesium on the solid surface may be heterocoagulaton of the pseudocolloidal form of Cs [13]. This mechanism will be discussed later. The exchangeable adsorption of ions on the metal oxides occurs in the presence of the ion type contaminations. On the surface of the oxide, beside the adsorption of the cation according to reaction 6, the substitution of the contamination for to the Cs takes place. = C t s + C s + ~-

-Cs s+Ct +

(8)

356 As far as the concentration of the cation on the surface of the oxide does not change with pH, the adsorption, according to reaction 8, is independent on the pH of the solution. On the other hand, because in the exchange reaction the H § ions may take part, then a small pH dependence of the Cs sorption can be observed. The investigation of caesium sorption on the titanium, aluminum and silicon oxides, performed by Hakem et al., revealed that the increase of the concentration of the electrolyte lowers the adsorption of the Cs-137 or 1-131 [14]. The pH dependence of the sorption of these radionuclides is typical for the adsorption of the ions on oxides. However, the higher adsorption of the cation at pH>pHpzc and anion at pH>pHpzc was observed. This behavior suggests that ion exchange process has a vivid share in the ion adsorption on the surface of the oxides. The authors of discussed papers characterize applied oxides by mentioning the size of the particles and specific surface, without telling about the existence of the ion contamination. Ionic impurities of metal oxide may have influence on the mechanism of the ion adsorption, especially from very diluted solutions-10-Smole/dm 3. The investigations made by Kosmulski et al. showed that porous glass, containing on its surface borsodium phase, is good adsorbent for Cs-137115-17]. The adsorption of this isotope is promoted by alkaline pH and low ionic strength. Some adsorption of the Cs-137 was observed on the silica gel [15]. Appropriately prepared four component glasses, Vycor-type, showed the good adsorption of Cs-137 [18-19], also not only in the alkaline pH, as it happened for three component glasses. Although the examination of the adsorption on the porous glasses focused on the obtaining the adsorbent for the removing the Cs radioisotopes from water, the achieved results showed that the presence of ions or the ion exchangeable layer on the surface of the oxide increases the adsorption of monovalent ions from the solution. Ion exchange character of the Cs adsorption on the soil sample that mainly consists of the sand was observed by Shenber and Johanson [20]. The adsorption of multivalent ions or monovalent hydrolyzable ions is specific adsorption. Because of the valence of the ion, more than one adsorption site may be occupied. The adsorption of hydrated form may go through dissociation of the hydrogen cation from the hydroxyl group of the adsorbed complexes as well as from the surface hydroxyl group. Because the adsorption of the metal cations on the surface hydroxyl group goes with dissociation of H § then the adsorption of cations in the some range abruptly increases with pH. This effect is called the edge of adsorption. The parameters that characterize the edge of adsorption [22], are explained in Figure 1. The specific adsorption may lead to the formation of inner or outersphere complexes [23,24]. As an innersphere complex is treated surface compound where the cation is directly connected with oxygen from the surface of the metal oxide (Figure 2a). The outersphere complex is formed when the adsorbed cation maintains the hydrated layer of water, Figure 2b. From the pH dependence on the adsorption, one cannot conclude, whether the inner or outersphere complex is formed.

357 ApHlo.9o%

20.0 --

6.4

9 -- ,.L

_--

o~ 15.0 - -

-- 6.0

E "O

_o O

E E

E N

10.0

0 e.0 0

<

e--

III tn(

_

"10

~+

!r

--

--

5.6

--

5.2

No-

a ) = d___Me

5.0-I

, 0.0

~

~-i

PHso% I 4

I

I 5

'

I

'

6

I

'

7

I

'

I

8

9

,

i

4.8

10

pH

Figure 1. Adsorption of Zn(II) (cricles) and concentration of Zn ions (triangles) in TiO2(Rutile)/electrolyte system as a function of pH (data from [21 ]). pHs0o/o- pH of the solution with 50% cation adsorption, April0-90 - the pH range where adsorption changes from 10 to 90%, dpMe/dpH - the parameter that shows how the activity of the cation must be changed with the changes of pH to leave the cation adsorption on the same level.

a)

b) (~

(Do 9 o ~I

0

o @

o~

o ~

o @

o

@

o @ ~t

o 9

o@

o ~

o9

0 ' e

000 0o~

k, ;

Hq H

o o

~I

o@

( D o @

o@

o@

sp ip

sp

IHP

Figure 2. Inner- (a) and outer (b) sphere complexes at the metal oxide surface, sp - surface plane, ip - innersphere complexes plane, IHP - Inner Helmholtz Plane.

358 The structure of the complex may be estimated by the spectroscopic measurements, as was proposed by Robertson and Leckie [22]. Especially, very useful are the spectroscopy methods that allow in situ investigations, i.e. Fourier Transformation Infrared Spectroscopy (FTIR) electron or proton resonance spectroscopy, M6sbauer spectroscopy or X-ray adsorption spectroscopy (XAS) [24]. These methods allow also to decide other mechanisms such as polynuclear surface complex formation or surface precipitation. The polynuclear complexes or clusters of the new phase may form at higher concentrations of the adsorbing ion. According to James and Healy, the surface precipitation of the hydroxide occurs at lower concentration of the cation than in bulk. That is because of the lower value of the dielectric constant in the compact region of the edl, than in the solution [25]. The coverage of greater and greater surface of metal oxide by the new phase, with the increase of pH (increase of "adsorption"), changes the surface properties of the oxide. At high enough coverage degree, there may occur the change of the surface charge from the negative to positive one in CR2 point (charge reversal). The following increase of the pH, results in the succeeding reverse of the charge sign CR3, at pH characteristic for pzc of the adsorbing metal cation hydroxide. Similar effect, of the charge reversal during specific adsorption of hydrolyzable cation, was described by Schindler [26]. The specific adsorption of an anion on the surface of oxides goes through the exchange one or two hydroxyl groups for the anion: k(- SOH)+ L n- +-~ (- S)kL n-k + k O H -

(9)

The greatest adsorption of the anion is at low pH and decreases from certain value of pH with the increase of pH of the solution. The weak acid anions may adsorb on the metal oxide surface from the solutions of the low pH as a molecule of appropriate acid with liberation of the water molecule: k(_ SOH)+ LHn-m m ~

n-m-k + k H 2 0 (- S)k LH m-k

(~o)

At the increase of the pH the respective ion forms of acid may be adsorbed. This type of adsorption occurs for As, Cr(VI), Mo and V ions. The cation complexation constants, on the surface of the oxide, may be calculated from the pH dependence, with the method proposed by Schindler [26]. Recently, the complexation constants are found with application of the numerical optimization, based on the chosen model of the edl (DLM, TLM). They allow to fit the model parameters to the experimental data (FITEQL [27], HYDRAQL [28], SURFEQL [29], GRIFT [30], SURCOM [31]). Fitting of the edl models with numerical optimization procedures allows to regard the energetic heterogeneity of the adsorption sites on the surface of the oxide [32,33]. A survey of data, concerning specific adsorption of cations on the metal oxides, was presented by Schindler [26], Schindler and Stumm [34], Kinniburgh [35] and Huang [36]. Some values for the complexation reaction

359 constants of the important for the environment heavy metals are presented in Table 1 and 2. Table 1 Negative logarithm of the apparent stability constants of surface complexes of heavy metal ions with one surface site Pb(II)

SiO2 5.1 [36]

Cd(II)

6.1

Hg(II) Co(II) Cu(II) H3AsO4 H2AsO4AsO43H3AsO3

[36]

5.52 [34]

FeOOH 3.8 [36] 4.65 [37] 4.9 [36] 0.47 [37] 7.76 [37] -0.46 [37] 2.89 [36] 29.31 [37] 23.51 [37] 10.58 [37] 5.41 [37]

A12Os 2.2 [35]

2.1 [33]

TiO2 0.2 [36] 0.44 [26] 3.2 [36] 3.32 [26] 4.3 [34] 1.43 [34]

Table 2 Negative logarithm of the apparent stability constants of surface complexes of heavy metal ions with two surface sites fl2'ii,,) Cu Cd Pb Co

SiO2 11.19 [34] 14.2 [34] 10.68 [34]

FeOOH 1.7 [38] 4 [38] 1.6 [38]

A12Os 7.0 [34] 8.1 [34]

TiO2 5.04 [34] 9.00 [34] 1.95 [34] 10.6 [34]

Other cation presence, for example alkaline metal cations, may change the adsorption of multivalent cations, by changing their activity in aqueous solution and competitive adsorption on the same surface sites. Because the adsorption of monovalent and multivalent ions on the surface of the oxides goes through the hydrogen ion exchange in hydroxyl groups, the competition of the adsorption may occur. For the adsorption that gives innersphere complexes, such competition may not occur. That is despite the reactions of hydrolyzable and background electrolyte cations, as the process is ruled by different mechanism. Moreover, the adsorption with innersphere complex formation is characterized by far greater adsorption IS OS constant. The relation I~Me(II) > > K.Mr > > I~Na can be observed [24]. Cations, which adsorb specifically with the formation of innersphere complex, reveal a shift of the adsorption edge (adsorption-pH dependence) for different concentrations of the

360 electrolyte solution. For example adsorption of Ba 2§ on TiO2 (anatase) [21], is depicted in Figure 3. In the case of innersphere complex, the adsorption edges for respective ion strength covers. This plot is demonstrated by the the adsorption of Cd on TiO2 (rutile) Figure 4. The adsorption of heavy metals on the metal oxides is the localized type adsorption and its isotherm (as a function of the concentration in the solution) is described by modified Langmuir equation. This modification includes interaction with surface potential: F,~/ FmaxKMe [H+ ~1 [MeZ+ ]* exp - RT FMe

(11)

=

1 + KMe [H+ ~1 [MeZ+ ]* exp(- ~-~/ For the settled pH value, the surface potential ~ maintains constant, so the components of the equation may be included, with ZMe, into adsorption constant Kad s = KMe

$[H+] -I *exp -

. Then the equation (11)is the same as Langmuir's

one. 0.05 - -

0.04 - -

r

E

"13 0 m o

E

0.03 - -

~l m

0

E o Q.

0.02

L

o (/) "13 <

0.01

--

0.00

' 3

I 4

'

I 5

'

I 6

'

I 7

'

I 8

'

I 9

'

I 10

'

I 11

pH Figure 3. Adsorption of Ba2§ ions at the TiO2(anatase)/solution of NaC1 as a function of pH, rectangles 0.001 mole/dm 3; triangles 0.01 mole/dm 3, circles 0.1 mole/dm 3 of NaC1, (data from [21 ]).

361

20.0

~E

15.0

+t% 10.0

O

o

<~

5.0

0.0

i

3

4

i

5

I

6

7

i

8

i

9

I

10

I

11

pH

Figure 4. Adsorption of C d 2+ ions at the T'O2~Rutile)/solution] ~ of NaC1 as a function of pH, rectangles 0.001 mole/dm3; triangles 0.01 mole/dm 3, circles 0.1 mole/dm 3 ( data from [21]).

Surface properties of soil are not uniform because of its composition, presence of different components and properties of crystal faces of each component. About this, for everyone component of the solid phase, one should use a separate set of the constants of the isotherm equation. Sposito showed, that it is possible to obtain the approximate Bemmelen-Freundlich isotherm, by using logarithmic-normal distribution of adsorption constants of the Langmuir isotherm [5]. In the case of very low concentrations of the cation, the adsorption may be described by Henry's equation. It may be applied thus for the adsorption of radionuclides having a short half-life time. At the relatively high concentrations, beside adsorption, the precipitation processes may take places. In the presence of the solid, for the cation concentrations higher or equal to the solubility product of the metal hydroxide, the adsorption isotherm does not break, that reflects the precipitation of the hydroxide of other phase but the isotherm changes the shape, which suggests the increase of the adsorption and is characteristic for the multicore complex formation or surface precipitation [24]. In the aqueous environment, there are some ligands beside adsorbing cations, which may form complexes with these cations. Then, in hydration sphere one or more water molecules may be substituted by a ligand. Depending on the complex type, the adsorbing complexing cation may be bind directly with the surface (A type complexes)

362

n - S O H + MeL~ + ~

(- SO- )nMe(L)I z-n)+ + nil+

(12)

or they may adsorb through a ligand (B type complexes) n - S O H + L M e z+ + nH20 +_~ (- SOH~)n LMe(n-z)+ + n O H -

(13)

In A-type complexes (reaction 12), as far as the formed complex does not adsorb or adsorb more weakly than the metal cation, the presence of L-ligand results in lowering of the metal adsorption. The possibility of the complex adsorption through a ligand, (for B type complexes) ease cation adsorption at lower pH values because the adsorption runs according to the mechanism characteristic for anions. Such influence of the ligand presence on the cation adsorption was observed for Ag§ 2 system. In the natural water environment, the following anions exist as complex anions: CI-, SO42-, HPO42-, F-, OH- and COa 2. Monovalent anions may influence the cation adsorption, according to reaction 13, whereas multivalent ones may increase or decrease the adsorption as well. The important role, as a complexing anion, is played by carbonate ions. Actinides and another elements which form with these ions the relatively stable complexes. Unfortunately, in the bibliography there is no data concerning cation adsorption in the presence of carbonates. Among many anions, potentially complexing heavy metal cations, present in a gmole amount in the soil many organic acids [39]. Their role was discussed in details by Hartera and Naidu [39]. 3.

CATION A D S O R P T I O N ON CLAY M I N E R A L S

Beside oxides, clay minerals, products of weathering of rocks, are important component of soil. They reveal lamellar structure, consisting of tetraheders sheet of XO4 ,,t" (X = Si, A1, Fe) and octaheders sheet of XO6 ,,o" (X = A1, Fe, Mg). Moreover, some places may be occupied by other cations of small size. Depending on the number of layers, the phyllosilicates may be divided into two, three and four layer types. Among these types in the soil are present kaolinite as twolayers ,,t-o" (Figure 5) and montmorylonite (vermiculite), smectite or mica, illite as threelayers ,,t-o-t" (Figure 6 )[39]. Generally, ideal phyllosilicates do not exist. The Si(IV) atoms in tetrahedrons may be substituted for Al(III) and AI(III) or Fe(III) atoms in octahedrons for Fe(II) or Mg(II). This substitution produces the resulting negative charge of phyllosilicate lattice, that is compensated by Na § K § or Ca 2§ metal cations, located in a position around ,,t-o-t" or ,,o-t" layers [40]. In comparison to the other minerals of this type, kaolinite has the relatively low specific surface (5-40m2/g) and low cation exchange capacity (CEC)[2]. The

363 threelayer minerals, containing water molecules or ions between ,,t-o-t" layers have greater surface for example: illites 100-200m2/g, vermiculites 300-500m2/g or smectites 700-800m2/g. The CEC value is slightly different and may be arranged in a following order illite < smectite < vermiculite [2].

o

0

o

o

,,~si o

~-

o

Si

0

0

A[

o,

O

sO

8i

8i

/

0

o-

8i

O O

\

0

9

OO AI

o .

H H

~

H

Figure 5. The model of the kaolinite structure.

t

O

t Figure 6. The model of structure of the three layer phyllosilicate.

364 According to Sposito, the electric charge on the mineral surface may form because of the isomorphic substitution of crystalline lattice atoms, or the reaction of the surface functional groups. The first one is called permanent charge[5]. For the hydrated oxides or twolayer phyllosilicates (kaolinite) is lower than 0.02 mole/kg, whereas for multilayer phyllosilicates (illite, smectite, vermiculite) the permanent charge is a hundred times higher. Because of acid-base reactions of the soil surface with H § or O H - i o n s the net proton charge is formed (determined as ~H). Innersphere and outersphere complexes also participate in the formation of the electric charge [41]. On the kaolinite, one can distinguish three kinds of the layers exposed to the surface: -the ,,o"-layer of the hydroxyl group connected with aluminum atoms, -siloxane ,,t"-layer, and edge plate of the ,,t-o" layer where =A1OH as well as --SiOH groups can exist. For the substitution of one Si atom by one atom of A1 in the siloxane layer, ,,t", the (=AI-O-Si-) group is formed, donated with negative charge. As it was previously mentioned, this charge is compensated by alkaline or alkaline earth cations. According to Stumm, the hydroxyl groups on the edge layer are characterized by constants pKs = 6.3 and pK~ 2 8.7, like hydroxyl group of A12Oa, =

whereas octaehedron display properties similar to gibsite pK~,~ ~-4. The surface charge comes from both kinds of groups, and pHp~c=7.5 [23]. Considering the solubility of kaolinite and the size of the permanent charge, Sposito found the surface charge density, point of zero net proton charge (p.z.n.p.c.) and point of zero net charge (p.z.n.c.) [42-45]. The determined value p.z.n.c.=3.5 is closer to iep value from electrophoretic measurements of kaolinite than 7.5. This position of p.z.n.c. suggests that the contribution of the hydroxyl group located on the edge in the formation of the charge may be smaller or the groups are more acid. On the basic of the surface charge density and zeta potential measurements, using the FITEQL program, D u e t al. found the ionization constants of surface groups of illite [46]. Because in the examined pH range the illite surface was negatively charged only pKa2 was determined. Good fitting was obtained using the model of energy homogeneous surface for constant pK=4.12-4.23 or for the surface of two kinds group. One of stronger acidity of pKa2(I)=4.17-4.44 and second type weaker acidity of pKa2(n)=6.35-7.74. Specific and nonspecific adsorption of monovalent cations (Cs § Rb+), in the face of high concentration of the ions in the soil solution, will be negligibly small, similarly to metal oxide/ aqueous electrolyte solution system. XPS and NMR examinations, also sorption of Cs § on the kaolinite, showed that considerable amount of caesium adsorb almost completely between the ,,t-o-t" layers, and only 1% of Cs adsorbs on the surface. The adsorption of weekly hydrated Cs § or Rb § ions in the region between layers is stronger in comparison with better hydrolyzed Li § or Na § ions, as these ions give lower potential formation of the edl [47]. The adsorption of cations of the higher radius (hydrated) promotes the swelling of the silt [3]. That explains the observed sequence of monovalent cation adsorption for silt type minerals [3,5,23]:

365 Li + < N a + < K + < R b + < C s +

(14)

The adsorption of the ions on the clay minerals is the exchange type process. The ions from the solution may substitute ions from interlayer for example: R - N a + +Csa+q ~

R-Cs + +Naa+q

(15)

where R - r e p r e s e n t s the phyllosilicate lattice with negative charge. The relative affinity of Cs § ions to the surface of phyllosilicate is described by the total selectivity coefficient: ZcsmNa Kc(Cs_~Na) = ZNamCs

(16)

where: Z- express the phyllosilicate fraction covered by ,,i"-ion (in relation to CEC), mi- molar concentration of ,,i"-ion (Cs § or Na § Though the equation for the selectivity coefficient resemble the equation for the reaction constant of the ion exchange, these coefficients are not thermodynamic constants. That is because the activity of ions in the crystalline lattice of the phyllosilicates are not known [23]. The adsorption of Cs § on the phyllosilicates is a complex process that runs through the succeeding stages. The adsorption of Cs § on K-illite and Ca-illite is well described by three-box model. Here the adsorption of the solution runs through two independent reactions followed by the third one, irreversible [48]. These mechanisms explain partial irreversibility of the Cs adsorption process on the phyllosilicates. It was proved that the adsorption on the illite covered with Ca ions (Ca-illite) goes better that on the mineral covered by potassium ions (K-illite). On the former also more Cs § adsorb in the irreversible way than on the latter. For this course of the reaction on both minerals, the higher distance between layers ,,t-o-t" in Ca-illite is responsible, that favors faster migration in the space between the layers. This migration is also responsible for the higher irreversible adsorption of Cs § on Ca-illite [49]. For divalent cation adsorption, the reaction of substitution runs through exchange of Me 2+ for two monovalent anions for example: 2 R - N a + +Me2q

~_ R2-Me 2+ +2Naa+q

(17)

and selectivity coefficient will be equal: ZMe m2Na Kc(Me-~Na) = Z~qam Me

(18)

366 The selectivity coefficients of the alkaline earth cation adsorption on the montmorylonite and vermiculite arrange in the following sequence [35]: Mg 2+
(19)

The investigations of the Sr 2§ and Ba 2§ adsorption on the kaolinite, illite and bentonite, covered earlier by K, Ca and AI(III) ions, proved that the process is well described by Freundlich or Dubinin-Radushkievich isotherms. The adsorptionconcentration dependence is poorly fit by Langmuir isotherm here. The adsorption of the ions is higher on the minerals covered K § than Ca 2§ or AI(III) [50]. For the adsorption of these ions the poor dependence from pH was noticed. The adsorption of heavy metal ions, such as Cd(II), Ni(II), Cd(II), Pb(II) is a more complicated process. Depending on the pH and concentration of the solution, beside the adsorption also the surface or bulk precipitation may take place. Cation adsorption on the silt type minerals may occur on hydroxyl groups of the layers exposed to the solution, on edges of the ,,t-o-t" layers (,,t-o" for kaolinite). It can run also by exchange with cations from the space between layers. It is considered that the siloxane layer (-Si-O-Si-) is not active towards the multivalent cations adsorption [35]. Schindler et al. found that the adsorption of Cd(II), Cu(II), and Pb(II) on kaolinite from aqueous solutions occurs on two types of sites, one of week acid character and another one on surface hydroxyl groups connected with aluminum [51]. The adsorption on week acid groups is ion exchange adsorption type and goes at low pH and ionic strength values. Higher values of these parameters favor adsorption on hydroxyl groups, resulting in the innersphere formation. The adsorption of Cu(II) and Pb(II) cations as a function of pH is similar for oxides/electrolyte systems, in certain pH range one can observe the increase of adsorption defined as the edge of adsorption. Differently, the adsorption of Cd(II) at increasing part of adsorption vs pH curve shows two distinct edges, one below and another above pH=6.5. Angove and coworkers, from the results of the Cd(II) adsorption and potentiometric titrations, found out the adsorption for pH_<4 is the exchange type adsorption on the permanent, negatively charged sites of the siloxane layer. This process is characterized by inconsiderable stoichiometry of a proton exchange (0.2) [52]. The adsorption on the hydroxyl groups takes place on the surface of the octahedron layer (A1) with characteristic, for alumina stoichiometry of the proton exchange. The -SiOH groups are in this adsorption of minor importance. In the similar way runs the adsorption of Co(II) in the system smectite/electrolyte solution [53]. Low values of pH and ionic strength promote the adsorption on permanent centers of the charge with formation of outersphere complexes, whereas increase of NaC1 concentration results in removing Co(II) from adsorption sites of permanent charge and favor multinuclear complex formation and surface precipitation. The size of these complexes increases with the increase of pH. XAS(X-ray Absorption Spectroscopy) investigations showed that distance Co-Co between atoms is smaller than in Co(OH)2. That may prove the adaptation of the

367 forming complexes to the crystalline lattice of the phyllosilicate. The smaller distances between adsorbed cations than between respective hydroxides were also observed for Ni(II) on many silty minerals [54]. In present paper, authors tend to opinion, that the smaller distance between adsorbed cations results from the formation of a mixing phase of the nickel-aluminum hydroxide. To the Pb(II) adsorption on the kaolinite data Majone et al. fitted different adsorption models. They stated that good fitting of the adsorption of Pb(II) as a function of pH and ionic strength of the electrolyte data is obtained using the threelayer model of the edl (TLM) characterized by two types of adsorption sites and their continuous distribution [55]. The adsorption of UO22§ on smectite runs in similar way to cation adsorption and is characterized by typical for them the adsorption edge [56]. The increase of Na § and Ca 2§ ions concentration in solution gives a shift of the edge towards higher pH values. In this system the adsorption model is very complicated because of many different ion forms of UO22§ and a solubility of smectite. In the investigations on the adsorption of Am-241 on the montmorylonite (the swelling mineral), illite (nonswelling) no significant difference was noticed, regarding their swelling behavior. That proves the actinides adsorption only on the surface of the minerals [57]. The adsorption of cations on the surface of silts may be described by the edl models characteristic for oxides. They should consider also the substitution adsorption processes on the interlayer surfaces of the silts, effects connected with energetic heterogeneity of the surface and dissolution of the solid. 4.

CATION A D S O R P T I O N ON C A R B O N A T E S

Calcium carbonates, beside oxides and silts are the main parts of soil components. Beside carbonates, recognized as minerals, other oxides may display surface properties of carbonates because of CO2 adsorption, though the structure of the bulk phase is oxide type. In this group may be included for example iron oxides [58]. A influence of carbon dioxide was also observed with the silt surface. Calcite and dolomite are the most popular carbonate minerals presented in soils, beside them siderite FeCO3 and rhodocrysite MnCO~ may be present. Because carbonates dissolve in aqueous solutions the determination of surface charge by potentiometric titration is difficult. The additional problem is caused by partial carbon dioxide pressure that influences the balance of the system. The behavior of the closed and opened to atmosphere systems is quite different. It is assumed that for the presence of anion and cation in the crystalline lattice of carbonates, on their surface the following reaction led to charge accumulation: -CO3H~=~

-CO~+H +

-C03H + Me 2+

~:~

-CO3Me+ + H +

(20) (21)

368 - M e O H ~ ~:t

(22)

-MeOH + H +

4:t -MeO-+ H + - MeOH+ CO2 ~ -MeCO~+H+

(23)

- MeOH

(24)

The charge density on the surface of carbonates is proportional to the algebraic sum of the concentrations of the respective forms bearing the charge [59], Go

: F{[-CO3H~]+[-C03Me+]+[-MeOH~1-[-MeO-]- [-MeCOg]- [-C03]}(25)

An electric charge on the surface of the carbonate, as of oxides, forms the electric double layer on the interface of the mineral. The presence of edl implies the occurrence of specific and nonspecific adsorption at the interface. Caesium ions may also adsorb on carbonates but because of higher concentrations of other cations of alkaline metals in the soil solution their sorption is probably minute. Divalent cations of heavy metals Co(II), Pb(II), Ba(II), Sr(II) show inclination to the formation of isomorphic crystals with Ca, Fe and Mn carbonates. Then, beside the adsorption in the edl a substitution adsorption may take place between Me 2§ ions from the solution and e.g. Ca e§ ions from a crystal lattice of the carbonate. This adsorption was observed for FeCOa/Mn 2+ aqueous solution and CaCO3/Cd 2+ solution systems [60,61]. In both, the adsorption process is more complicated the substitution adsorption is the first quick stage, then a diffusion to the hydrated layer of the carbonate proceeds and finally diffusion into the solid. (CaCO3XCaCO3 * H20)Cd 2+"

(26)

(CaCO3XCdCO3*H20)s+Ca2+

(27)

(CaCO3XCaC03 * H20) s + Cd 2+ ~ (CaCO3 XCaCO3* H20)Cd s2+~:t

(CaCO 3 XCdC03 * H20)s + c a 2+ ve

((Ca, c a ) c o 3 XCdC03 * H20)s + Ca 2+

(28)

In Davie's et al. opinion, the plot of Cd 2§ adsorption versus pH confirms the above mechanism of the process, the adsorption lowers with the increase of pH. Despite the presence of hydroxyl groups on the surface, the substitution adsorption process cannot run with hydrogen ion liberation. In these condition the increase of pH is accompanied with the increase of the adsorption, for example metal oxide/Me 2§ aqueous solutions. The observed decrease of adsorption Davies et al. explain by the exchange for Ca 2§ ions [61]. After all, this interpretation of adsorption versus pH dependence seems difficult to accept because in the examined pH range (pH6-8), the concentration of Ca 2§ ions decreases [62]. When assuming the Cd-Ca ion exchange mechanism, the increase of the adsorption with the increase of pH (decrease of H § concentration) should be observed. The increase of Me 2§ (Me=Cd, Zn, Mn, Co, Ni, Ba, Sr) adsorption on calcite was observed by Zachara et al. [63]. In their paper the following sequence of the cation adsorption on calcite was found:

369 Cd 2+ > Zn 2+ > Mn 2+ > Ni 2+ >> Ba = Sr

(29)

From the experiments on the divalent cation adsorption, the character of their substitution adsorption with Ca 2§ ions was proved [63, 64]. Additionally the following dependence was observed: the higher hydrating energy of the cation, the easier its desorption. Beside the Me 2§ ion adsorption, also the dissolution and recrystallization processes occur in carbonate/electrolyte solution system. These processes may be responsible for formation of solid solutions of carbonates (Me 2§ ion transportation into the solid phase). The influence of the recrystallization on the diminution of the cation from the solution is visible at longer adsorption time, and may be interpreted as "diffusion" into the solid phase. During recrystallization, the cations from the surface are covered and then the process lead to incorporation of adsorbed ion into crystal lattice of a solid phase which is one of the mechanisms typical for the system consists of cocrystallised micro- and macro constituents [13, 61, 65]. The description of Cd(II) adsorption process on CaC03, which considers recrystallization of the solid was proposed by Das and Van der Weijden [66]. 5.

A D S O R B I N G A N D C O M P L E X I N G P R O P E R T I E S OF ORGANIC S U B S T A N C E S OF S O I L

Considering their origin, organic substances presented in soil may be divided into two groups. The first one includes all substances, result from the natural, biological processes that happen in the environment, the second one contain all substances introduced by man and his industrial activity [3]. In the first group there are substances of small and big molecular weight such as acids, amines and aminoacids. The most important carboxylic acids are: oxalic, formic, citric, acetic, succinic, malonic, maleic, aconitic and fumaric. Their concentration in cultivated soils is lower than in forest ones [39]. The group of substances of the high molecular weight includes lignins, celluloses, simple proteins and products of their degradation. All these organic substances present in soil show various properties and different solubility in water solutions. An amount of the organic substance dissolved in water is defined as a dissolved organic carbon (DOC) and measure in mg of carbon/dm 3. To the group of organic substances of the anthropogenic origin may be included detergents, sulfonic acids, whitening agents, polymers, solvents, fuels and so on. Some of them dissolve well in water and form complexes or precipitates with heavy metals. Organic substances presented in soil and called "humic substances" do not have well-determined chemical constitution. They are formed mainly because of biological degradation of lignins, proteins and carbohydrates (mainly cellulose). Substances of smaller molecular weight, formed at the beginning and bonding

370 together gives humic macromolecules in the end [5]. For the different substances may react in this way the final particle may has the variety of functional groups: COOH, -OH (phenolic and alcoholic)=CO, -COH, -NH, -SOH, N in heterocyclic orimides and amides [5]. The average molecular weight of fulvic acid is 670 a.m.u. and molecule contains six carboxyl and five phenolic groups [4]. The average humic acid particle is bigger and may reach 25 000 a.m.u.. The humic substances reveal properties similar to polyelectrolyte gels because of their three-dimensional structure. The dissociation of carboxyl or phenolic group forms the electric potential around the humus molecule or their gel phase (for the insoluble humus). To neutralize this charge, some cations from the electrolyte accumulate at humus molecule [67]. The analysis of the H § affinity to the functional groups of humic acid, based on the model of continuous adsorption sites (energetic heterogeneity), showed two peaks on the distribution curve. One for carboxyl groups of p K H Int ~ 4 and second characteristic for phenol groups pKHInt -8-9 [68]. Ephraim and coworkers believe that humic acids behave in intermediate manner between simple electrolyte and polyelectrolyte; as oligoelectrolytes with discrete distribution of acidity constants of functional groups [69]. Marinsky et al., proposed the method for the calculation of ionization constants of carboxyl groups and ionization constants of metal cation complexation including chelate complexes [70]. They proved that fulvic acid contains three types of carboxyl groups (I,II, III), characterized by different dissociation constants (pZaI - 1.2, pZaII - 3.4, and pZaIII - 4.2) and acidic alcohol (enol) groups of the constant pKaIv =5.7. For metal complexation, there is possibility to obtain four unidenate and four chelatic species[70]. The organic substances of small molecular weight, show complexing behavior not only with metal cations in solutions but also may complex these cations that are in solid phase, for example A1 or Fe. This action promotes the solubility of oxides or other minerals [71]. The solubility of fulvic acid complexes with A1 or Fe cations, depends on their mutual ratio (metal cation/fulvic acid). For the ratio equal to one the complex is soluble. For higher ratio values (3-6) the solubility decreases [4]. Acid properties of humic substances may be learned by potentiometric titrations [23,72]. From these data the distribution of pK constants of humus functional groups may be found [73]. The complexing properties of humic and fulvic acids are the subject of many papers dealing with cation adsorption and the references can be find in many reviews [4,5,38,71]. Harter and Naidu presented values of complexation constants for some heavy metal cations: Co, Cd, Cu, Fe, Mn, Ni, Pb, Zn, by ligands existing in aqueous solutions of soil [39]. The metal cation adsorption on humic acid particles is complex process because of polyfunctional character of acid group, its polydispersity and existence in dissolved and colloidal form [38]. For simplification the adsorption reaction is assumed to run on quasiparticle with groups of the acid character: aSHn(aq or s)+ pMe 2+ + qL1- + xH + + y O H - ~-where: 5=2p-x-y-lq and b = n ' a ,

SaMep(OH)xLSq(aq or s) + bH+ (30)

371 Above equation may be reduced to simpler form regarding that humic acid is dissociated in aqueous solution [38]: a L H - + Me 2+ ~-

MeL2a(l-a) + a H +

(3~)

This equation is characterized by the constant KH: KH = [MeL2(1-a) ][H +

(32)

[LH- ~ [Me 2+ ] It can be noticed from the above equation, that the metal adsorption on humic acid should increase with the increase of pH. This behavior was observed for Pb(II), Cu(II), Cd(II) and Ca(II) [38]. The organic substances with groups of acidic (or alkaline) character are active towards molecules of opposite character or to surface groups having alkaline(or acidic) properties. Molecules of humic substances may adsorb on the soil minerals by the interaction of their functional groups with mineral surface. The presence of aliphatic chain segments or aromatic rings enables humus particles to the disperse interaction. Sposito distinguished following mechanisms of the organic substances of soil: - cation exchange, - protonation, - anion exchange, - water bridge formation, - cation bridges, - ligand exchange, - hydrogen bonding, - Van der Waals interaction. Humic acid adsorbs on the oxides and on kaolinite in the way characteristic for anions, the adsorption decreases with increase of pH [23,71]. Because of the size of particle and n u m b e r of functional groups the adsorption may run only for the part of molecule, such as for other macromolecules. Thus, h u m u s particles may immobilize the nonionic particles (organophosphates halogen derivatives of hydrocarbons) alkaline type organic particles or metal cations [4,23]. Humus particles compete with organic acid particles of small molecular weight in adsorption and complexation reactions [71]. The adsorption isotherm of humic acid on the kaolinite does not show the tendency to reach plateau, characteristic to complete coverage, so the humic acid adsorption is multilayer type. The reasons for this are nonpolar interactions of hydrophobic segments of acid molecules [74]. It was observed t h a t humic acid adsorbs in lower degree t h a n fulvic at the same concentration because greater particles of humic acids cover up neighbouring

372 adsorption sites. The adsorption of fulvic acid, which particles are smaller, is a Langmuir type (monolayer adsorption) [74]. The adsorption of Cu(II) on the kaolinite, covered previously by humic or fulvic acid showed that: - adsorption affinity of Cu(II) ions to humic acid is stronger than to fulvic acid - adsorption affinity of Cu(II) ions does not depend on the molecular weight of humic acid, and is always the same. - Cu(II) and H § ions adsorb completely on the same adsorption sites. In the soil/electrolyte system coexist: a solid inorganic phase, humic substances and solution that contains, among others, the metal cations. Model investigations of such systems are done mainly for metal oxide (or phyllosilicate) - humic acid electrolyte solution [75-85]. The presence of DOC changes the mechanism of cation adsorption because of cation complexation. On the one hand, the complexation leads to lowering the cation adsorption at higher pH values, according to the mechanism of reaction 12. On the other hand, the cation adsorption on the solid takes place, according to reaction 13, through the adsorption of ligands. For humic acid, because of the existence of many different functional groups, the adsorption may run according to both mechanisms (reactions 12 and 13). This effect was observed for the Cu(II) adsorption on A1203 and presence of humic acid [75]. Fulvic acid lowers the Zn(II) adsorption on goetite and hydroarargite by complexing the metal cations in the solution (reaction 11) and has no influence on this cation adsorption on SiO2 [85]. The adsorption of rare earth elements, in the presence of humic acid, is complex process and differs much from the solutions without this acid [76-80]. This adsorption is connected with humic acid complex formation with Eu(III) ions in the solution, adsorption of these complexes also the adsorption of ionic form of Eu(III) on the adsorbed humic acid on the oxide or on the phyllosilicate. Similar effect of humic acid on the adsorption, was observed for ions of V, Ag and other rare earth elements [85]. Beside discussed group of ions of rare earth elements (group A), Takahashi et al. distinguished three groups of ions of elements. They differ in their adsorption behavior in the environment humic acid/ dispersed oxides (phyllosilicates) [85]. Group B - (Mn,Zn, Co, Be, Sr, Ba, Fe, Cr) - humic acid causes desorption of these cations in neutral pH. Group C - (Ru, Rh, Ir, Pt, Ga, Zr, Hf) - humic acid does not have influence on the cation adsorption but earlier adsorbed acid molecules limit Hf and Zr ions adsorption through screening the adsorption sites. Group D - (Rb, As, Se, Te) - the presence of humic acid does not have any influence on the adsorption of these ions. The description of adsorption process in the metal cation - humic acid - mineral system is far more complicated than for the system without acid. Till now, there are some methods used for the description of adsorption in systems containing humic (fulvic) acid; LOGA J.C.M. de Wit et al. [86], NICA model (by Bendetti et al.) [87]. One can also adopt adsorption model that considers energetic heterogeneity of adsorption sites [33].

373 Another review dealing with actinides complexation by humic substances was presented by Maulin et al. [88]. In this paper some complexation constants of selected lanthanides and actinides were given. The complexed forms of actinides with humic acid dominate in solutions of pH<7 (sometimes 8) at concentrations 0.1 mg/dm 3. Whereas the presence of other cations (AI(III), Ca(II)) may change the contribution of h u m u s complexes of actinides in aqueous phase. To obtain the adequate opinion of the existing forms in the system, the calculations may include the concentrations of anions and cations in the aqueous environment of soil.

Q

P R O C E S S E S OF FORMATION, T R A N S P O R T A T I O N AND A D S O R P T I O N OF COLLOIDS IN SOIL SYTEMS

Heavy metal cations and actinides may often precipitate under the soil solution conditions. The average concentration of anions that form insoluble sediments in ground or surface water is often sufficient to form carbonates, hydroxides, sulfides, phosphates, fluorides or chlorides with respective cations. If in discussed system some colloids are formed, then processes characteristic for real solutions cannot explain the ion transportation in natural environment. Obtained as a results of radioactive isotope precipitation, fine sediments (insoluble salts) are called real colloids. It was mentioned earlier that radioisotopes under specific conditions may adsorb on fine dispersed oxides or hydroxides of colloid size. Then, the behavior of this system will not be the same as for real solutions, but rather as the colloid systems. Isotopes, adsorbed on colloid matrix form pseudocolloid [3]. Beside two mentioned colloid systems, Kim distinguishes in natural systems the third kind aqua colloids. They are formed as a result of succeeding reactions: dissolution of the mineral, hydrolysis of obtained product, polinucleation and colloid formation [89]. Depending on the solid contacted with water, the aqueous solution can contain even more than 100 ppm of colloids. Usually high concentrations of aqua colloids exist in water that contact with humus. Actinides form insoluble sediments with hydroxides, carbonates, sulfates, phosphates and fluorides [90]. There are no data concerning the solubility of actinides with silicate anion, though uranium forms with this anion several insoluble minerals. Because in natural conditions in ground water the hydrocarbon and carbonate anions play the dominant role, in the reactions of precipitation (in the 0.01 mole/dm 3 solution of Na § pH=7, pC02 =3.5) depending on the oxidation state of the actinides, the compounds presented in Table 3 may precipitate [90]. The behavior of actinides in natural environment was described in many papers reviewed by Lieser [3,89], Silva and Nitshe [90], Kim [91,92], Chopin and Stout [93], Newton and Sulivan [94], Larsen et al. [95] and Tanaka et al. [96]. The behavior of colloids or pseudocolloids in natural soil and water systems is a subject of the investigations in the aspect of the transportation not only the toxic substances but also gaining precious minerals or elements, for example Au [97]. The examination of the behavior of Pu in alluvial sediments of Los Alamos depository revealed that Pu loaded on the colloid, translocates few times slower

374 Table 3 Solid phase and solubilites of actinides[90] Oxidation state of An

Solid phase

+3 +4 +5 +6

AnOHC03 AnO2 NaAnO2CO3 AnO2(OH)2*H20

Solubility of An [mol/dm 3] 10 .7 10 -10 10 .5 10 .6

than tritium but more than thousand times faster then it was predicted by two or three phase model [6]. The model investigations of colloid transportation through columns packed with quartz, showed that the migration rate of bigger particles, for example latex, is greater than that of smaller ones [98, 99]. To describe the latex transportation in a column, the model similar to dynamic chromatography was proposed. The transportation in quartz packed column is treated as a particle transportation through capillary. The average rate of particles (v), depends on equivalent radius of capillary Ro, the rate of the liquid (velocity profile fluid) Vr, rate of particle Rp and energy of interaction between particle and capillary (packed quartz) W [99].

Vtr),exE Wr)lrr (v)=

kT

J

(33)

~:~ -RP e x p l - :T(r)]rdr The energy of interaction consists of Van der Waals disperse interaction WVDWand electrostatic interaction of edl WDL W = WD1 + WVDw

(34)

For RpfRo << 1 the interactions between particle and package of the column (quartz grains) may be treated as sphere - plate interaction. Then, the electrostatic interaction will vary with particle potentials ~gl and the plate potential ~2 WD1 (h)= 16~ .Rp.

tanh ewe//.tanh e~2/.exp -Kh 4kT J \ 4kT J

where: h- distance between plate and sphere e- elementary charge ~- dielectric constant K-reciprocal Debaye length K = ~

1000e2N Av -~ ~i ziM 2 i

(35)

375 NAv- Avogadro's number z - valence of the ion M - molar concentration. The above equation couples the influence of electric potentials on the surface of particles and ionic strength of the electrolyte on the electrostatic force interactions. An increase of ionic strength results in drop of WDL by the increase of • parameter. For spherical particle - plate system the Van der Waals force interactions are described by following equation: WVDW (h) = -

+~ + 21n x+l x+l

(36)

where x=h/2R~ A- Hamaker constant. Hamaker constants for pure components of discussed systems are available in many monographies dealing with stability of dispersed systems [100]. For real systems, for example latex-quartz interactions in water, the H a m a k e r constant have to be calculated from pure component data.

Wv w: (A, 11-

X

(37)

where: All, A22 H a m a k e r constants for solids (1- latex and 2 - quartz), A33 Hamaker constant of the liquid medium (water). Velocity of particle migration in the porous systems, calculated in this way, in the opinion of Nagasaka et al., may well predict the behavior of colloids in geological systems[99]. This model contains many simplifications: it assumes a small size of colloid particles in comparison to the packing of the column, does not take into account precipitation on and liberation particles from collectors, heterogeneity of the package, and dynamics of processes and at last variation of chemical and physical conditions. In bibliography, there are known also the more developed models of porous collectors and the particles transportation [101-104]. However, despite many models of the colloid particle - collector interactions, which may be successfully adopted to well-defined systems, their application in such natural systems as soil still does not give the precise description of the transportation process. Comprehensive, excellent review of the state of present knowledge on the colloid transportation in the natural environment was presented by Ryan and Elimelech [6].

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