Precipitation of metal-humate complexes

J ,n.m nuU Ch, m Vol. 43, pp, 1219-12-'22,1981 Prir~ltrd ~n (ircil tlritai~

O022-1902,tgl'(Y3121t)qt4502 (W)[O Pergamon Pres:; Ltd

PRECIPITATION OF METAL-HUMATE COMPLEXES GORDON K. PAGENKOPF* and C. WHITWORTH Department of Chemistry, Montana State University, Bozeman, MT 59717, U.S.A.

(Received 19 May 1980; received[or publication 19 September 1980) Abstract--Concentrations of magnesium found in sea water are capable of precipitating cadimum humate from water solution. Speciation calculations and kinetic considerations indicate that dissociation of the cadmium complexes occurs during the precipitation process. For a sequence of divalent metal cations, lead is the most effective precipitator of humic acid, whereas magnesium is least effective. The order is lead > copper > cadmium :, barium > zinc > nickel > calcium > strontium > magnesium. INTRODUCTION The transport of trace metals by natural water systems is facilitated by a sizeable number of processes. One of these is complexation by naturally occurring chelating agents [1 ]. With few exceptions, these naturally occurring chelating agents, humic acid and fulvic acid, have not been well-characterized. These materials result from degradation of plant material, and they are present in virtually all waters. Complexation studies with transition and trace metals have provided stability constants, conditional stability constants, and mass action quotients for complex formation in several systems[2-16]. Most of these constants were obtained using ion selective electrodes, ion exchange, or electrochemical techniques. The conditional stability constants increase in value as the pH of the water increases. In addition, the higher molecular weight ligands are observed to bind more than one me~al ion per molecule[ll, 13]. The interactions are believed to primarily involve oxygen donors; however, sufficient nitrogen is usually present to provide additional potential coordination sites. We have studied the reactions between coal-derived humates and many divalent metal ions. A sequence of complexes is formed that progressively increases the metal to ligand ratio. This results in ligand charge neutralization, and at a high ratio of metal to ligand ( > 10 to 1) the metal-humates precipitate. Specifically, alkaline earth metals have been utilized to precipitate cadmium humate. The results of the study indicate that fresh water soluble complexes will precipitate upon mixing with waters that contain calcium and magnesium at concentrations similar to that found in sea water.

observing the intensity of 90° scattered light. The mercury vapor light source and a 2.5cm cell were utilized. Each measurement was made relative to a solid standard to account for light source and detector drift. All light-scattering experiments were conducted at pH 7.0. The precipitated humic acid complexes were separated by centrifugation at 15 K for 15 rain, using a Beckman JA-20 rotor. The obtained precipitate was dissolved with concentrated nitric acid at 100°C, and the metal content was subsequently determined by atomic absorption spectrophotometry. Total organic carbon analyses were done with an Oceanography International total carbon system. RESULTS AND DISCUSSION Humic acid solutions were titrated with magnesium, calcium, strontium, barium, nickel, copper, zinc, cadmium and lead. Previous studies involving cadmium indicate that humic acid is capable of complexing me)re than one metal ion per ligand, and in fact, cadmium ions are successively bound until a sizeable fraction of :the negative groups are neutralized[l l, 13]. The sequential binding of metal ions progressively reduces the net charge of the complexes, which eventually results in precipitation. The onset of precipitation was monitored by recording the intensity of scattered light. The titration of 50.0 ml of 20.5 mg/l HA with 1.00 x 10-2 M Cu 2+ at pH 7 is shown in Fig. 1. The total concentration of copper required to cause 50% change in scattered light intensity is 3.6 x 10 5 M (0.182 ml). Using

EXPERIMENTAL Coal humic acid (HA) was obtained from eastern Montana sub-bituminous coal. The acid was extracted by sodium hydroxide in the presence of dioxygen, with isolation and purification similar to that described by Malcolm[17]. A detailed discussion of the isolation procedure and characterization has been presented elsewhere [13]. Fhus humic acid is similar in chemical composition to those isolated from other coals and soils[18], and contains 0.3% ash. All solutions were prepared from reagent grade chemicals using doubly-distilled CO2 free water. The metal stock solutions were standardized using EDTA and an appropriate indicator, pH electrodes were standarized using commercially prepared pH 4 and pH 7 buffers. All titrations were carried out at 25.0°C, I : 0.10 M KNO3 under nitrogen. The titrant was added with a calibrated Gilmont micrometer syringe buret. Humic acid solutions were prepared by weighing approx. 10 mg HA followed by dissolution with 50.0 ml of 1.00N HNO3. The solution was then diluted to 500.0 mi with distilled water. The formation of the insoluble complexes was monitored by *Author to whom correspondence should be addressed. ~lN~v . ~; No ~ H

1219

4 -----~

/

/ c

/

.,7. R

i

/

/ ,

o

7 O0

, Ol

!

,

,

I

02

03

rnt 0.0~0 M Cu2+

Fig. 1. Dependence of scattered light intensity on added copper. [HA]T = 20.5 mg/1, pH = 7.0, I = 0.10 M KNOs, initial volume = 50.0ml.

G. K. PAGENKOPFand C. WHITWORTH

1220

the previously determined average gram formula weight of 6761[13], the total HA concentration is calculated to be 3.0× 10-SM. Ion selective electrode measurements with copper indicate that the conditional stability constants are greater than those for cadmium. As a consequence, the free copper concentration is too low to monitor by electrode. As the total copper concentration increases, the free copper concentration increases, but since the stoichiometry is unknown, conditional stability constants cannot be calculated. The data in Fig. 1 indicate that precipitate formation starts when the total copper concentration is 2×10-5M (--0.1ml), and is complete when the concentration equals 6 × 10-SM ( 0.3 ml). A solution that had a total Cu(II) concentration of 3.5 x 10-5 M was centrifuged and the filtrate and precipitate analyzed for Cu. From Fig. 1 this amount of copper should result in approx. 50% precipitation. Analysis provided 1.4 x 10-5 M Cu in the filtrate, and the precipitate contained an amount equivalent to 1.6× 10-5 M. By mass balance 14% was lost. With a total HA concentration of 3.0 × 10 6 M, a Cu/HA ratio of near 10:1 is indicated for the precipitate. Total carbon analysis provided 14.1mg/1 carbon in the HA solution before copper addition, and 4.5 mg/1 carbon after precipitation. For precipitation coincident with the addition of one additional Cu(II) ion CuxHA-" + C a 2+

Ki

Cux+,HA -'+2

(1)

the equilibrium constant, K, for the reaction will be greater than 104 M-1 at pH = 7.0. The amount of divalent metal ion needed to cause 50% change in scattered light intensity for all metals studied is summarized in Table 1. All of the metals caused essentially the same change in scattered light intensity during the precipitation process; thus, the amount which caused a 50% change in scattered light intensity was selected for comparison purposes. The most effective are lead and copper, with magnesium being the least effective. Comparison of the stability constants for complexation of these metals by ligands such as acetate and formate[19] provides a qualitative agreement between total metal required for 50% change in scattered light intensity and the strength of complexes formed with oxygen donor ligands. The strongest complexes require the smallest concentration. The formation of hydroxide complexes will influence the order for the metals that partially hydrolyze at pH 7. Table 1. Metal concentration needed to cause 50% change in scattered light intensity, pH = 7.0, HA = 3.0 × 10-6 M Metal

Concentration,

Lead

2.9 X 10-5

Copper

3.6 X 10-5

Cadmium

2.2 X 10-4

Barium

2.7 X 10-4

Zinc

4.7 X 10-4

Nickel

9.6 X 10-4

Calcium

2.9 X 10-3

Strontium

3.8 X 10-3

Magnesium

7.3 X 10-3

The observation that many metal ions are capable of precipitating humic acid suggests that mixed metal complexes would also be insoluble. Specifically, we were interested in the question of whether cadmium would be lost from the humate complex before the humate could be precipitated by an alkaline earth metal. To study this, cadmium was mixed with humic acid at a concentration ratio that would insure partial complexation of cadmium. Magnesium was then added and the resulting precipitate was digested and analyzed for cadmium. The data are presented in Table 2, first five entries. These results indicate that 4-8% of the total cadmium is precipitated. The other entries in Table 2 contain sufficient cadmium to cause partial precipitationn of Cd:HA~s~. In these cases, the amount of cadmium precipitated increases with an increase in total cadmium; however, the percentage precipitated remains fairly constant at about 3%. The last three entries indicate that Ca, Sr and Ba are comparable in the precipitation of the cadmium complexes. The equilibrium species distribution for CdHA, CdzHA, Cd3HA and CdaHA in the five solutions of lowest total added cadmium is presented in Table 3. The distribution was obtained using the program COMICS [20] and neglected the formation of any cadmium hydroxide species. The conditional stability constant for the formation of CdHA, Cd2HA, and Cd3HA at oH 7 from Cd2+ and HA are 1059, 101°6 and 1015"7 respectively[13]. The value for /~'1 at pH 7 is 102°2, and a value for /~'1 at pH 7.5 is 10259121]. The uncertainty in the first three values is estimated to be -+0.2 log units. It is larger than this for/3'~ and/3'~. Within experimental uncertainty, the successive stability constants decrease as the number of metals bound increases. This is consistent with the observation that a total cadmium concentration of 2.2 × 10 -4 M is required to cause a 50% change in scattered light intensity. The species distribution calculations for the lowest CdT study show that 49% of the metal is complexed, 6.55 × 10 8moles, with the predominate species being CdHA. Addition of magnesium results in the precipitation of 1.0× 10-Smoles of cadmium, which indicates that a sizable fraction, at least 85%, of the complexed cadimum dissociates during the precipitation process. As the total cadmium concentration increases, the species distribution shifts toward higher order species, i.e. CdaHA, see Table 3. Correspondingly, a larger amount of cadmium is precipitated by the addition of alkaline earth metals. In general, the rates of water exchange for these divalent metal ions are large[23] and thus the metal distribution in the precipitated humate may be kinetically or thermodynamically controlled. The cadmium exchange rate may be estimated using the conditional stability constants and a predicted value for the formation rate constant. For the equilibrium kl

CdxHA " + Cd2+ = Cd:,+~HA .+2 kr

the forward rate constant, k t, is equal to the following: k~ = k n2o × Kos where k-H2o is the rate of water loss from cadmium and Kos is the outer sphere association constant. The water loss value for Cd(II) is 108.2sec -~ [23, 24]. Values for the

Precipitation of metal-humate complexes

1:221

Table 2. Cadmium--humate precipitation by alkaline earth metal ions. pH = 7.0 humic acid = 3.0 × 10-~ M ml Metal

Molarity

Mg2+

5.4 X 10-2

Total Volume Total Cd, moles

CdT, M

Cd ppt. moles

25.1

1.3 X 10 -7

5.3 X 10.6

1.0 X 10 -8

25.1

2.2 X 10-7

8.9 X 10.6

1,8 X 10 .8

25.1

9.0 X 10 -7

3.6 X 10-5

3.2 X 10 -8

25.3

1.8 X 10 -6

7.1 X 10-5

1.3 X 10-7

25.4

2.7 X 10 -6

1.0t5 K 10 -4

!.0 X 10 -7

25.5

3.6 X 10 -6

1.41 < 10 -4

1.3 X 10-7

25.6

5.3 X 10 -6

2.07 ~ 10.4

2.1 X 10-7

25.8

7.1 X 10 -6

2.75 ~ 10-4

2.4 X 10-7

26.0

9.0 X 10 -6

3.46 X 10-4

2.6 X 10 -7 2.5 X 10-7

Mg2+

1.2 X 10 -2

57.0

1.4 X 10-5

2.5 X 10-4

Ca2+

5.6 X 10-3

53.0

"

2.6 X 10-4

3.8 X 10 -7

Sr 2+

5.6 X 10-2

53.0

u

2.6 X 10-4

4.3 X 10-7

Ba2+

4.8 X 10.3

52.5

"

2.7 X 10 .4

4.0 X 10 .7

Table 3. Cadmium-humate species distribution, pH = 7.0 105Cd~

106Cd2+

I06CdHA

I07Cd2HA

10.7_Cd_3 HA

I08Cd4HA

.53

2.69

1.83

2.47

.84

.71

.89

5,12

1.79

4.59

2.95

4.78

3.60

26.8

0.28

3.78

7.10

60.3

O. 04

I . 30

9.83

187.

98.4

O. Ol

O. 62

7.39

222.

I0.6

a)

12.8

108.

HAT : 3.0 X 10 -6 M

outer sphere association constants may be calculated from first principles. For reactions involving a divalent cation and a neutral ligand, the value is near 0.15 M ', with a monovalent anion the value is near 1 M " , and for a divalent anion the value is approx. 40M '[24]. The effective charge on humic acid isn't known, but it is probably at least - 2 for each cadmium. Using Ko+ = 40 M L ks = 108.,- x 40 = 6.3 x 109 M ' sec '. The conditional stability constnats vary from 1059 to 10+5; thus, using the relationship k, = Keq/kt the values for k, range from 103.9 to l053 sec i. Rate constant of this magnitude indicate that the half-life for dissociation of the complexes is less than a millisecond. As a consequence, the stoichiometry of the humic acid complex that precipitates is probably thermodynamically and not kinetically controlled. And correspondingly, the metals that form the most stable complexes will be preferentially precipitated, provided that the rates of complex formation and dissociation are rapid compared to the rate of precipitation. This should be the case for all of the metals included in this study. In summary, humic acid can be precipitated by divalent cations. The mechanism presumably involves charge neutralization that is accompanied by loss of hydration energy. Magnesium and calcium present in sea water should be very effective precipitators of humates,

and provide a route for the transfer of trace metals from the solution phase to the solid phase. Acknowledgements--This research was funded in part by the

U.S. Environmental Protection Agency, Grant No. R803727, Cincinnati, Ohio. Appreciation is extended by C. W. for ~L National Science Foundation Energy-Related Graduate Traineeship. REFERENCES

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G.K. PAGENKOPF and C. WHITWORTH

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18. M. Schnitzer and S. U. Kahn. Humic Substances in the Environment, pp. 325, Marcel Dekker, New York (1972). 19. L. G. Sillen and A. E. Martell, Stability Constants of Metal1on Complexes, 2nd Edn. The Chemical Society, London (1964). 20. D. D. Perrin and I. G. Sayce, Talanta 14, 833 (1967). 21. C. Whitworth, Dept. of Chemistry, Montana State University, Bozeman, MT 59717, June 1978 (unpublished observation). 22. M. Eigen, Angew. Chem. 3, 1 (1964). 23. M. Eigen and R. G. Wilkins, Adv. Chem. Series 49, 55 (1%5). 24. R. G. Wilkins. Accounts Chem. Res. 3, 408 (1970).