Coordination Chemistry - American Chemical Society

Chapter 33

Humic and Hydrous Oxide Ligands in Soil and Natural Water Metal-Ion Complexation

Downloaded by COLUMBIA UNIV on January 21, 2018 | http://pubs.acs.org Publication Date: November 4, 1994 | doi: 10.1021/bk-1994-0565.ch033

Cooper H. Langford Department of Chemistry, University of Calgary, 2500 University Drive NW, Calgary, Alberta T2N 1N4, Canada Werner complexation and the coordination theory play an important role in understanding the chemical processes in natural systems in water and soil. However, the ligands are more complicated than those usually encountered in the laboratory because they are commonly polymeric, colloidal -scale, heterogeneous, and polyelectrolyte. The ideas required for understanding simple ligand equilibria and kinetics must be supplemented with a sensitivity to the effects of a polyfunctional distribution of sites that is subject to electrostatic effects arising from the polyelectrolyte character and conformational and aggregational equilibria, which may be coupled to the extent of metal ion coverage. This paper develops these points using soil humic substances and hydrous metal oxides as examples.

Natural Ligand Donor Sites Although some of the complexing sites for metals that exist in natural waters and soil solutions are located on molecular ligands such as amino acids, simple oxo anions, and hydroxyacids, the main and most persistent ligand sites are on larger particles. In the few-nm domain we find organic ligands such as the water-soluble low-molecular-weight (ca. 10 ) humics known as fulvic acids and inorganic species with cation-binding capability such as A l ( O H ) and the somewhat larger F e ( O H ) polymers. Ascending the size scale we encounter the humic acids (larger cousins of the fulvic acids), polysaccharides, and small hydrous oxide colloids. Near the usual boundary for an experimental "definition" of dissolved matter (the 0.45-μπι filter) lie typical F e O O H , M n 0 , and clay particles as well as cellular debris and metal sulfides. Organic compounds adsorbed on inorganic particles function as 3

7 +

1 3

x

y

2

0097-6156/94/0565-0404$08.00/0 © 1994 American Chemical Society Kauffman; Coordination Chemistry ACS Symposium Series; American Chemical Society: Washington, DC, 1994.

3 2


33.

LANGFORD

Ligands in Soil and Water Systems

405

members of the size class of the adsorbing substrate. Thus this coordination chemistry is focused on the domain between the molecular-level ligands and "suspended solids." If the physical scale of the ligands introduces complexities which were not of concern to Werner, the nature of the donor sites on these intermediate scale ligands adds no complexities. Structures 1, 2, and 3 could stand as representatives of the vast majority of the complexing sites for metal ions, (M), with many Ο and a few Ν donor atoms. Sulfide sites and simple halide ions are the other donors of lesser importance. Consequently, we need to anticipate nothing in the bonding of the ligand to the metal that was not exhibited in the complexes known to Alfred Werner. This soil/water system chemistry is the chemistry of classical complexes, mainly between metal ions and hard base ligands. Ligand Heterogeneity A fulvic acid of number average molecular weight 1000 can yield as many as 30 different substituted mononuclear phenol-carboxylate structures in a mild oxidative degradation which accounts for 30% of the mass of the sample. Because only a small fraction of these structures would be needed to make a polymer of M W = 1000, it is clear that the fulvic acid is a complex mixture, yet the aliphatic components have not so far been considered. It may be asked if efforts should not concentrate on the resolution of the mixture, regardless of how difficult that task may be. However, the task in understanding natural system chemistry would then require modeling the interactions that occur in the mixture, which is, in my opinion, more difficult than studying them directly. Buffle(7) has introduced an extremely helpful concept, that of a homologous group of ligands. He recognizes that there are polymeric, polyelectrolytic ligands which vary in details of molecular structure but have common features in the basic metal-ligand bonding. Such homologous groups will not be fully characterized at the molecular level, but it is useful to divide the study of coordination chemistry into the study of the various homologous groupings and to characterize the behavior of the contributing homologous groups. Examples of groups of homologous ligands are the fulvic acids, humic acids, natural polysaccharides, nonhomodisperse Fe(III) or Al(III) hydrous oxides, and clay particles. In all of the homologous groups the three factors which complicate the treatment of metal ion equilibria and kinetics are: (1) the polyfunctional character and distribution of molecular type of donor site; (2) change in conformation on metal-ion binding, coiling, aggregation, and gel formation; and (3) polyelectrolyte character changes in electrostatic effects on metal-ion binding. Homologous ligand groups can be subdivided into three types, which differ mainly in the roles of these three factors: type I, dissolved polyfunctional complexant; type II, monofunctionalpolyelectrolyte (permeable

Kauffman; Coordination Chemistry ACS Symposium Series; American Chemical Society: Washington, DC, 1994.

406


COORDINATION CHEMISTRY

R & S are generalized groups



33.

LANGFORD


407

gel); and type III, monofunctional polyelectrolyte (surface reactant). Table I summarizes the differences among these types and gives examples of each. Consider the specific cases of a fulvic acid, a simple polyfunctional complexant, and the surface of a hydrous ferric oxide particle, a typical monofunctional polyelectrolyte. In the first of these cases, the variety of ligand structures in the mixture will require recognition of a distribution of binding constants to describe the system. In addition, both conformational change in the polymer and variable electrostatic effects will make these component binding constants themselves appear to be functions of the degree of site coverage by the metal ion. The most useful approximation has been to assume that the distribution of component binding constants is a continuous function of site coverage. In the hydrous oxide surface case, the treatment can be somewhat simpler. A l l of the complexing sites are variants of the surface hydroxyl type. Interest is usually directed to the binding of a trace metal, present in low concentration compared to the number of surface sites. This may be called the geochemical limit because it describes common natural concentration ratios. Only the most strongly complexing sites are likely to be important, and it is often a very good approximation to assume a single surface ligand binding constant for a metal and to correct only for polyelectrolyte effects. Kinetically, the situation is not simple for either of the ligand types. Especially, the number of kinetically distinguishable sites seems to exceed the number distinguishable at equilibrium in the hydrous oxide surface case. This may well be a consequence of sites of similar character thermodynamically being kinetically differentiated by the length of the diffusion path into the particle interior, which must be traversed by a metal ion in the course of binding or dissociation. Equilibrium and Kinetics of Binding to Hydrous Oxides A typical study observes the uptake of trace metal ions on a dispersion of a hydrous oxide, yielding the percentage of metal adsorbed (% ads) as a function of p H . The p H range over which the variable % ads changes substantially is usually narrow, leading to the term "adsorption edge." Figure 1 shows data from Schindler(2), one of the pioneers of the interpretation, for uptake of Cu(II), Pb(II), and Cd(II) on hydrated rutile ( T i 0 ) . The lines show a fit with the theory which assumes a single intrinsic binding constant,*K , for reaction (equation 1) and a single intrinsic constant, */3 , for reaction (equation 2). Corrections have been applied for polyelectrolyte character based on the Gouy-Chapman-Stern theory. 2

lint

2int

Surf-OH + M 2Surf-OH + M

z +

z +

2

+

+

= Surf-OM^ *> + H = (Surf-0) M< > + 2 H z2

+

(1) +

2

(2)

Figure 2 shows the persuasive correlation between the intrinsic constants used in the treatment of surface complexation and the stability constants of simple hydroxo complexes measured in homogeneous solution.


408


Table I Characteristics of the Main Homologous Complexing Groups

i

n


Dissolved polyfunctional conplexant

m Monofunctional polyelectrolyte (permeable gel)

Monofu nctional polyel ectrolyte (surfac e reaction)

Schematic structures

Λ

: °: 0

1 *

fulvic compounds PROM AROM

little hydrated Fe, Al, K> oxides ] c

a y s

Examples peptides strongly hydrated proteins metal hydroxides humic compounds polysaccharides cell walls metal oxides and clays covered by NOM i polyfunctional character b change in conformation c polyelectrolyte character

strong

intermediate

intermediate-weak

aggregation

formation coagulation spreading/coiling

important

important

important

The horizontal positioning of examples depends on their behavior, e.g., the intermediate position of proteins indicates that in this case properties I and II are both important. O.N.S. = coordinating atoms of complexing sites. (Réf. 1).



33.

LANGFORD

409


Figure 1. Adsorption of some divalent metal ions on Τ Ί 0 (rutile). For each metal ion there is an interval of 1-2 pH units where the extent of adsorption rises from zero to almost 100%. The solid lines are calculated with the constants calculated for single-site model. (Reproduced with permission from Chimia. Copyright 1976 Chimia Abodienst.) 2

5

Figure 2. Correlation of stability constants log * K C^onnt)) complexes at amorphous silica with stability constants log *Kj (*p ) of hydroxo complexes. The solid line represents the equation log *K (*^ (int)) ~ ~ + 0.62 log * K , (*β ). (Reproduced with permission from Oester. Chem. Z. Copyright 1985 Verlag Lorenz.) o

f

s

u

r

f

a

c

e

1 ( i n t )

2

5

=

1(int)

2

2


0 0 9


410


A somewhat different experiment is illustrated in Figure 3(5). Here the adsorption isotherm for the uptake of Cu(II) ion on and into colloids of hydrous ferric oxide is reproduced. The Cu(II) ions were present in the acid solution prior to the neutralization steps leading to formation of the Fe(III) colloids. The isotherm is simple and can easily be fitted within experimental error to a single binding constant approximately consistent with those obtained in adsorption edge experiments despite the somewhat different uptake mechanism. However, a kinetic experiment in which the Cu(II) was stripped from the hydrous iron oxide with a Cu(II)-selective ligand required a minimum of four kinetically distinct Cu(II) types to account for the data. Figure 4(4) shows the trend in the concentrations of the four kinetic components as a function of the Fe:Cu ratio in the initial solution. The largest component (top curve in the figure) is the labile component comprised of free Cu(II) and Cu(II) bound to surface sites and accessible for reaction in less than 3 s. (The kinetic technique will be explained in the section on ligand exchange kinetics.) Equilibrium of Binding to Humic Substances The study of binding to humic substances has been greatly aided by the availability of ion-exchanged samples rendered into the protonated form with all other cations being rendered negligible. It is then possible to cany out stoichiometrically well-defined experiments on the titration of these ligands with metal ions. A classic example is illustrated by the three titrations shown in Figure 5(5). The titration of the Armdale soil fulvic acid with Cu(II) was monitored with three different experimental probes. The filled circles represent the monitoring of the Cu(II) bound by means of the measurement of "free" Cu(II) with a Cu -ion-selective electrode (ISE). The open circles represent the monitoring of the quenching of the fulvic acid fluorescence at 465 nm. The triangles represent the monitoring of the increase of absorbance (and light scattering) at 465 nm as a result of Cu(II) complexation. The quenching of fluorescence, Q (see left axis), is smooth and suggestive of a relatively simple binding function. The fraction of stoichiometrically characterized sites covered as indicated by ISE, χ (left axis), increases monotonically but clearly indicates more than one binding constant. The change of optical absorbance, A (right axis), has two inflection points, strongly suggesting some physical changes in the ligand during Cu(II) uptake. The last point will be documented by light scattering data (Figure 7). The fluorescence data suggest that the fluorescing sites are a subset of the stoichiometric total, probably associated with stronger binding constants. A separate experiment using gel filtration chromatography has shown that the fluorescence is associated with the lower-molecular-weight fractions separated from the mixture in the gel filtration process. The ISE experiments were treated quantitatively on the assumption of a continuous distribution of binding constant. The results in the form of a plot of the free energy of 2+

0

4 6 5


33.


LANGFORD


^

411

20

ο 0

10

20

TOTAL C u

30 2 +

40

(10~

6

50

60

M)

Figure 3. Example of Fe-hox binding of Cu(II) determined at p H 6.8. Fitting of these data to conventional isotherms shows good adherence to Langmuir theory. (Reproduced with permission from reference 3. Copyright 1989 National Water Research Institute.)

100

Fe:Cu RATIO Figure 4. Results of applying kinetic speciation analysis to the equilibrium binding of Cu(II) by Fe-hox at pH 6.0. ·, X component 1; 0, component 2; • , component 3; Δ, not recovered, component 4. A minimum of four kinetically distinguishable components are necessary to describe the system. (Reproduced from reference 4. Copyright 1993 American Chemical Society.)


412



Figure 5. Comparison of three experimental probes of metal binding to the fulvic acid, χ , the fraction of sites occupied by C u as calculated from ISE titration; Q, the relative quenching of F A fluorescence; A ^ , the absorbance by fulvic acid at 465 nm as a function of C u added. See text for guide to curves. (Reproduced from reference 5. Copyright 1981 American Chemical Society.) 2 +

2 +



33.


LANGFORD

413

binding vs. fraction of sites covered is shown in Figure 6. In this figure the left axis divides the free energy of binding into two terms, an intrinsic chemical term, AG°, and a term for the electrostatic effect expressed as a ratio of activity coefficients, Γ. The partition is conceptual only; it was not actually performed. The role of conformation and aggregation is spectacularly indicated by the light scattering data shown in Figure 7. In this figure the rise of right-angle Rayleigh scattering as a function of the coverage of sites by Cu(II), estimated from ISE data, is shown. A t all p H values the initial uptake of Cu(II) leads to little change in light scattering. A t higher coverage uptake of Cu(II) leads to a sharp increase in aggregation with a slope steep enough to suggest a cooperative process. It was proposed that the initial portion reflects the uptake of Cu at didentate sites. A s these become saturated, Cu is bound to monodentate sites. This allows the formation of bridged complexes of the type L - C u - L . However, if L is a polymer which can interact with another L by hydrogen bonding and/or donor-acceptor interactions, the bridged complex may have a closed form as shown by the sketch in the figure. The interacting ligands can form a pseudo-chelating ligand, and this may account for the apparent cooperativity. The interaction is promoted by bridging, and it stabilizes the bridged species. Kinetics of Ligand Exchange from Humic Substances As mentioned in the context of hydrous oxide surfaces, the kinetic behavior of naturally occurring complexes is important. In addition to the fact that kinetics may show a different partition of complexing sites from equilibrium experiments, there is the issue that biological uptake of metal ions is a dynamic process and that kinetics may control the effects of metal ions on biological systems. Because there are few sensitive methods to detect complexation reactions in progress, we adopted ligand exchange between the metal complexed with a humic substance and a complex of that metal with a spectroscopically sensitive ligand (R, below)(6). This is not only experimentally beneficial, but it also relates to the biological situation where a metal ion is often taken up via a surface complex with a cell membrane. Consider the complexes of Ni(II) with the acid-soluble soil humic fraction, fulvic acid ( F A ) as they react with a reagent R to form a spectrophotometrically detectable product. There is a set of competing reactions: 2+

Ni(aq) + R -» N i R F A N i + R -> N i R + F A FA -Ni + R NiR + F A r

2

FArNi + R

X

2

(3)

NiR + FAi.

The subscripted FAj represent the kinetically distinguishable groups of binding


414


Cu-FA

\

FA= 0-100 λ

pN « 6-0 NO KCI


\

15-0

• \

\

7 ο

Ε M

10*0

\

H•S

•\

+

\ ·

\

°β I 5-0

0

j 0-2

1 0-4

1 0-6

Figure 6. Results of titrations with C u at pH 6.00 shown as the free energy of binding derived from the Κ values. AG is the intrinsic term. Γ represents the electrostatic term. (Reproduced from reference 5. Copyright 1981 American Chemical Society.) 2 +

0


33.

LANGFORD

415


sites on the F A mixture. The first reaction, with free aqua Ni(II) is the fastest. Provided that a large excess of R is used, all reactions will be experimentally pseudo-first order. More fundamentally, the reactions may often (but not always) be described by the following mechanism: slow 2+

FA Ni

> FAj + Ni (aq) fast > NiR

r

2+


Ni (aq) + R

(4)

In this favorable case the measured first-order rate constant for each component will be the dissociation rate constant for the metal complex of the corresponding complexing site or site collection. With all reactions reduced to pseudo-first order, the experimental rate law with excess R takes the form: d[NiR]/dt =

A [ l - exp(-kit)] + X .

(5)

G>i

In equation 5 A represents the initial (time zero) concentration of the ith complex species expressed in units consistent with the measure of the product N i R , kj is the rate constant for the ith species, and X is a time-independent term which contains the spectroscopic blank and the contribution of any species which reacts too rapidly for the time scale of the experiment. The blank can be measured independently. The task is to fit the experimental signal-time function to equation 5, identifying a set of rate constants, k and initial species concentrations, A . Because the equation is nonlinear, this is not easy. The process has three steps(7). First, it is necessary to identify objectively a suitable minimum number of kinetic components, i . Second, the best nonlinear fit to the data must be obtained. Finally, the analysis must be tested for its chemical validity. This must be done to assure that the calculations have accomplished more than simply a set of numerical parameters which form no more than a way to archive the data. (We cannot overemphasize the importance of this last step. It is extremely easy to generate numerical analysis artifacts of little chemical interest.) The first step is now accomplished using a numerical approximation to the Laplace transform of the signal-time function(#). The numerical differentiation used in the transform always requires data smoothing. The second step can be accomplished using nonlinear regression on the unsmoothed data. The chemical validation is usually accomplished by studying the kinetics over a range of concentration and p H conditions of samples. The indication of validity is the stability of the rate constants and the evolution of the concentrations as expected from mass action considerations. Figure 8 shows the evolution of the four components identified in a study of the F A - Ni(II) system(9) as a function of F A : N i ratio in the samples at p H = 4. Q is free aqua Ni(II), C is the most labile complexes, and C and C o i

l9

o i

2


3

4

416


20 ,-

PH= 3.6 ο 6.0 • 8.0 Δ

15

ο

ο

σ

>

ιη 1

^

0

ο" Downloaded by COLUMBIA UNIV on January 21, 2018 | http://pubs.acs.org Publication Date: November 4, 1994 | doi: 10.1021/bk-1994-0565.ch033

Φ

C

Α—Μ—Β

ο 5

ο

L

pseudochelation

ο OÛ ο

_l 2.0

1.0

3.0

BOUND Cu

5.0

4.0 mMOL-9

Figure 7. The behavior of relative right-angle light scattering R ^R° at various p H values alues ;as bound C u increases. Bound C u is calculated from ISE titration. (Reproduced from reference 5. Copyright 1981 American Chemical Society.) 9

2 +

9Q

2 +

LiJ

3

4

5

6

7

10

FA:Ni RATIO Figure 8. Changes in relative contribution of initial components ( φ as a function of ligand-to-metal ratio following equilibration at pH 4.0. O , Cjî • , C ; V, C ; and 0 , C . (Reproduced from reference 9. Copyright 1987 American Chemical Society.) 2

3

4


33.

LANGFORD


417

represent less labile sites. A s expected if the least labile complexes are the strongest binding, their concentrations rise with increase of F A , the labile (weak) complexes increase but show signs of leveling off. Free Ni(II) declines. Rate constants for species 3 and 4 are approximately 2 χ 10" s" and 3 χ 10" s" , respectively, and are the only species whose lability may be too low for the dynamic bioavailability to be equal to that of free aqua Ni(II). 2

3

1

1


Conclusion The novel features introduced into coordination chemistry by study of the colloidal ligands which are critical in soils and natural waters are three. First, we must deal with a distribution of similar but distinct binding sites. Second, polyelectrolyte effects are prominent. Third, the ligands undergo conformational and aggregational equilibria which are influenced by metal ion uptake. Literature Cited 1. 2.

3. 4. 5. 6. 7. 8. 9.

Buffle, J. Complexation Reactions in Aquatic Systems: An Analytical Approach; Ellis Horwood Limited: Chichester, England, 1988. Schindler, P. E.; Stumm, W. In Aquatic Surface Chemistiy: Chemical Processes at the Particle-Water Interface; Stumm, W., Ed.; Wiley: New York, NY, 1987; Chapter 4. Gutzman, D. W.; Langford, C. H. Can. J. Water Pol. Res. and Control 1989, 23, 379-387. Gutzman, D. W.; Langford, C. H. Env. Sci. and Tech. 1993, 27, 13881393. Underdown, Α.; Gamble, D. S.; Langford, C. H . Anal. Chem. 1981, 53, 2139-2140. Langford, C. H.; Kay, R.; Quance, G. W.; Khan, T. R. Anal. Letts. 1977, 10, 1249. Langford, C. H.; Gutzman, D. W. Anal. Chim. Acta 1992, 256, 183-201. Olson, D. L.; Shuman, M . S. Anal. Chem. 1983, 55, 1103. Lavigne, J. Α.; Langford, C. H.; Mak, M . K. S. Anal. Chem. 1987, 59, 2616.

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Coordination Chemistry - American Chemical Society

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