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I '9 LIBRARY ,- MiChigan S *a'a University THE RELATIVE STABILITIES CF METAL-ORGANIC COMPLEXES IN RECENT LAKE SEDIMENT By Susanna Yetunde Sonaike A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of MASTER OF SCIENCE Department of Geology ABSTRACT THE RELATIVE STABILITIES OF METAL-ORGANIC COMPLEXES IN RECENT LAKE SEDIMENT By Susanna Yetunde Sonaike A series of potentiometric titrations was carried out to examine the effects of metal interaction with organic-rich sediment from Lake Lansing, Michigan. The objective was to determine the relative affinities of several metals for orga- nic-rich sediment with respoct to complexing and displacement. A constant amount of well—stirred sediment to which 50 ml of distilled water has been added was used for each titration. Known amounts of several environmentally—important, complex— forming metals were added, and a titration curve was determined for each using O.1N KOH as titrant. The metals used were: Cr+3, Fe+2, Fe+3, Co+2, Ni+2, Cu+2, Zn+2, Cd+2, Hg+2, Pb+2. Relative stabilities for each complex were computed using the method of differential precipitation. ACKNOWLEDGEMENTS Acknowledment is due to Dr.Charles Spooner who freely gave his advice and guidance at every stage of this investigation, and to Dick Kotch and Dane Williams who were always ready to lend a helping hand. I should also thank my husband, Femi, for his moral encouragement and physical help in compiling this work, and little Tayo-for making things a little less difficult during those critical months. ii Chapter Chapter Chapter Chapter Appendix TABLE OF CONTENTS Introduction Sample Collection Experimental Technique Theory Of Potentiometric Titrations Chelate Formation Results and Discussion Summary and Conclusion iii Page O\\O \lO‘x {Tl-T H Table 1: Table 2: Table 3: Table A: LIST OF TABLES Changes in pH and Equivalence Points Relative Stability of metal complexes of Lake Lansing sediment Summary of Relative Stabilities of Humic Acid fractions in four studies Confidence Level of titration points iv Pages 14 15 i (Appendix) Fig. Fig. Fig. Fig. Fig. LIST OF FIGURES Titration Of Sediment: in absence of metal ion Titration Of Sediment: in presence of SOOOmg of metal ion (Cr'H't, Cu , Fem ) Titration Of Sediment: in presepce of 5000mg of metal Cation ( Pb“; Co“, Ni ) Titration Of Sediment: iqfipresence of SOOOmg of metal Cation ( FeES Fe + Cu Titration Of Sediment: in presence of 5000mg of metal Cation ( Mn++ Zn++ Ca++) Titration Of Sediment: in presence of 5000mg of metal Cation (Cd++ Hg++ Hg+++ Cu++) Titration Of Sediment: Precipitation CHAPTER 1 Introduction What is the "fate" of metals diSposed of into lakes? In these days of profound environmental concern, the answer to this question is of major importance. Do metals disposed into lakes remain in solution on the overlying water or are they taken up by the organic sediment to form complexes? If the latter is the case, what can we say about the relative stabilities of these metal complexes? These are the two prime questions which this experiment attempts to answer. Radiotracers are a method that can be used in determining the competition of metals in sediment. However, there are a number of problems involved in using this method. This method requires a labo- ratory that is well—equipped in nuclear material. Thus extra precau- tion have to be taken in handling the radioactive material. By contrast, the potentiometric method, used in this study, offers a safer, less complex and equally reliable technique for measuring 5 the competition of metals in any natural environment. However, the Q (I) K . radiotracer method could be useful in providing plementary infor- mation on the competition kinetics of metals in natural environments. For this study, a series of potentiometric titrations were carried ced C out using sediment from Lake Lansing. Metal ions we-e introd 1‘ into the sediment before titration with 0.1N KCE. The aim was to deter— V mine whether or not chelation occurs between these metals an: the sediment. The relative stability of each metal complex was also 1 determined. According to Martell and Calvin(1952), the pH of a solution is directly affected by chelate formation since all chelating agents are also acids (or bases), and formation of the chelate is accompanied by the displacement of a proton from the acid, or by a decrease in the concentration of the base. The potentiometric method is a good technique for detecting chelate formation. Hoffman(1969) points out the importance of dissolved organic complexes in natural aqueous environments. In many such environment, especially in sediment pore waters, organic complexes probably act as the controlling factor for solution or precipitation of minerals. Organic complexes cause cation activities to behave as complicated functions of Eh, pH, and the total amount of dissolved organic matter. Presence of these complexing agents can cause the dissolu- tion of minerals whose cation concentrations is at a level of saturation. A knowledge of the relative order of stability of metal comple— xes is useful to the understanding of ore formation. Helgeson(1964) has shown that the deposition of sulphides in ore prOportions is a function of changes in the stabilities and degrees of formation of complexes in solution. Such deposition may be localized over a short depth range by dilution with cool ground water, or by its mixing with other hydrothermal solutions of different composition. Complexes of the non-ore—forming metals may be responsible for sulphide deposition or hydrothermal leaching through their changing degrees of formation in response to varying geothermal and composi- tional environments. Whether or not a metal ore would be deposited depends on the stability of the metal complex (ore-forming metals -3- are generally transported in the form of metal complexes), and on the type of chelating agent. Helgeson(196h) computed the order of stability for chloride complexes of the ore-forming metals at 25 degrees centigrade as: Cu++ 1— Zn++ 4 Pb++ 4 Ag++ A. Hg++ This proved to be the reverse of what was observed for the metal complexes of the organic sediments in this study. CHAPTER 2 SAMPLE COLLECTION Sediment sample was collected from the deepest part of Lake Lansing. This was done with a bottom dredge lowered from a small boat. The dredge was hand-lowered with a rcpe and tripped With a messenger. The contents of the dredge were placed in a plastic colle— ction bag which was then sealed and taken to the laboratory. About 250 ml of this sediment was put in a one—litre glass beaker. Dis— tilled water was added up to 1/h inch from the tOp of the beaker and covered with an airtight cover. This beaker was kept in a water- bath at a temperature of 200C for a week. EXPERIMENTAL TECHNIng' A small vial of about 7.5cc in volume was used to measure the amount of sediment used for each potentiometric titration. Samples were taken from the 1—litre beaker of Sediment. The sediment in the beaker was stirred for about ten minutes, until a uniform mixture was obtained. Several small vials were filled with the sediment and sealed. For each titration, the contents of one vial was emptied into a 100 ml beaker and 50cc of distilled water was added. This was again stirred thoroughly and then allowed to settle. O.1N KOH was prepared and used for the titrations. A 5ml burette was used. Procedure: An Orion Research pH meter with Beckman electrodes was used. The meter was standardized with a buffer solution of pH 10. The first titration of the sediment was done in the absence of metal cation. Each titration was carried out by adding a few drops of the titrant. at a time, stirring the sample, giving it time to settle, and recording the changes in pH (see Appendix). Subsequent titrations were carried out by adding 5,000mg of diffe- rent metal solution to each sample before titrating. Standard stock solution of the following metals were used: Cu++, Zn++, Cr+++, Co++, Ni++, Pb++, Hg++, Cd++, ++ +++ ++ Fe , Fe , and Mn . The titration curves for each metal are presented shortly. A sediment sample from the small vial was dried at 1050C and weighed: Weight of sediment = O.h519gm Volume of water added to sediment before each titration = 50ml Density of sediment = O.h519 = 0.009gm/cc Density of sediment at each titration = 0.009gm/cc. CHAPTER 3 THEORY OF POTENTIOMETRIC TITRATIONS Potentiometric titrations are concerned with the reactions of ions. Electrolytes may be classified as strong or weak. By the theory of Arrhenius, strong electrolytes in solution are in the main dissociated into ions and the degree of dissociation decreases with increasing concentration of the electrolyte. Consider a O.1Molar solution of a strong electrolyte BA. We have: [Bf]: [A‘] = 0.1 = 10'1 where, ’ + . . . B:] is concentration of cation and, [A2] represents concentration of anion Following in the tradition of Sorensen, we can eXpress the ion concentration as the ion exponent, PI, which is the negative loga- rithm of the ionic concentration. Thus: 1 P1 = —10 1 = g[ ] Iogrg] During a titration, the ion concentration changes over a considerable range and in a pattern elaborated upon by Kolthoff and Furman(1926) (Chapter 1). The Equivalence Potential: The potential of a metal electrode is a linear function of the metallic ion exponent in the solution. Therefore the change in potential during a titration is an indication of the change of the metallic-ion eXponent. At the theoretical end—point of a titration, there is, as a rule, a sudden change of the ion exponent which is indicated by a corres— ponding jump of the potential at the electrode. Thus the electrode can be considered a specific indicator for the corresponding metallic— ions in the solution; consequently, it is called an indicator electrode Titration Curves and Detection of Equivalence Point: The change of the pH of the solution and the variation in the potential of the electrode during titration, can be represented by a plot of pH versus volume of titrant added. This then will be a titration curve. The titration curves obtained under different conditions are shown in Figures 1-6. The maximum change in potential occurs at the equivalence point, or, if the titration curve is not symmetrical. at both sides of the point, very close to the equivalence point. This means that the equivalence point is located at the point where: < ‘H . . .. figg— reaches a max1mum, or its 2nd derlvatlve d d N E :1: Close to the equivalence point, we add the reagent drOp by drOp and note the readings when the potential remains constant with additional drops. CHELATE FORMATION The main purpose of the eXperiments was to find out whether or not complex or chelate formation takes place when metals are presen- ted into organic-rich sediment. The pH of a solution is directly affected by chelate formation since all chelating agents are also acids (or bases) and formation of the chelate is accompanied by the displacement of a proton from the chelate. The H+ liberated can be titrated with NaoH or KoH. pH Effect: According to Martell(1952), all metal chelates may be considered as formed by the displacement of one or more usually weak acidic protons of the chelating agent by a metal ion. Thus the addi- tion of glycine to a solution of cupric salt causes a drop in the pH due to the equilibrum: co —— 0\ NH —— CH + 2 2 Cu+2+ 2H NCHZCOO' —=- / Cu 3 ‘c‘CHZ—bfi \o—co The reaction may also be considered as the combination of the metal ion with the anion of the chelating agent, accompanied by a shift in the ionization equilibrum of the amino acid: 00—— 0 NH -———-— CH Cu+2+ H NCH C00" *1“ | \c / 2 2 (:I~12--——-1\:H2 o ---—— co The result is the same in both cases. Although it is difficult to fully and accurately define the orga- nic Species in sediments, we can safely assume that the complexing agents include humic acids, amino acids or folic acids. In that case, the reaction would be as eXplained above. In a qualitative sense therefore, we can conclude that chelate formation of metal ions with amino acids and with the acid form of other chelating agents results in pH drop(Nartell,1952). The greater the tendency for metals to combine with a given chelating agent, the greater the pH drop. In this eXperiment, the addition of metal ions to the sediment resul- ted in an appreciable drop in pH for some metals (fig.3) and a rise in pH for some others (figs 2, 5). 1 . TITRATION or SEDIMENT ”9" in absence of metal ion vol. of MN KOH anus.) _ A TITRATION 0F SEDIMENT in presence of 5000ug of Fig.2 :. " ' , metal. ion = a. 12 4..» _ I, fi 0 .2 vol. 'of on: «on Guns) A TITRATION or SEDIMENT in 'presence of 5000ug of metal Cation _ _ vol. of OJN KOH (mu .- 15 20 5 TITRATION OF SEDIMENT ‘_ Fig.4 in presence of 50001.19 of A metal Cation _ L. O ‘1: F: t CUE . / ' ‘ ml J .__ .5 vo|.10 of 50.1N20K0H25( 3» 30 TITRATION or SEDIMENT in presence of 5000 pg - of metal . Cation vol. of cm KOH (mist Fig.6 TITRATION OF SEDIMENT in presence of 5000 pg of metal Cation 1 vol. 2 of 301N_ 4KOH 5(mls.) Fig. ’7 -IJTRAIION .oF _ SEDIMENT: FREQ PITATION l — F. ‘ ' - s_-_.-.-_-:_-_—__-_._ Turbr' J -__— _-_— _- - - -_ .__, Cle d7, {17.3.2 :1 ’1: 323253”? 3-:{.—.—."—_§‘_'§;— ”mama We" gar-22? , _ "5301ng ' éé"ff€i._$ed1men+ .1263?“ ,3: :2) » ;Be+‘o're Th'rafl'on- After Ti+'ra+lan . CHAPTER h RESULTS AND DISCUSSION TABLE 1 Changes in pH and Equivalence Points (pH of pure sediment = 7.55) Metal pH change on pH at equivalence Cbmment adding metal point(precipi— to sediment tation begins) Divalent- transition 41* Mn++ 8.05 9.30 Fe++ 2.05 5.85 Fe (0H) ppt CO++ 2.15 5.20 pink pgt Ni++ 2.25 5.u0 blue ppt Cu++ 6.35 8.65 blue Cu(0H)2 ppt Zn++ 6.35 7.55 Zn(OH)2 ppt Trivalent— transition Cr+++ 8.05 8.80 yellow ppt Fe+++ 3.05 5.80 brownish not of Ee(OH3) Other divalent Pb++ 2.u5 5.35 ++ Cd 7.05 9.u5 Hg++ u.30 5.05 first inflection point 0a++ 3.50 o.u0 * No compleXing occurred with these metals \0 -10- Explanation of what happened during the titrations: Formation of chelate is accompanied by a drop in pH, due to the . . + . llberatlon of H lons. For example: M+2 + Ke-————)'MKe + 2H+ . , + When the hydrogen ions are titrated With a base (e.g. hOH), the H liberated is neutralized by the OH" from the base to form water. Therefore, when all the H+ are neutralized, that is at the equivalence point, the metals come out of the sediment to form hydroxides such as in: Cu(QH)2 ————— which gives blue ppt. 00 ———— pink ppt .++ light blue ppt Nl++ “_—‘ . Fe ———— yellOWlsh brown ppt Fe+++ ———— brownish ppt For those metals which do not form insoluble hydroxides there was no precipitate: the metal merely remained in solution (Fig.7). The equivalence point was extrapolated from the titration curve and also from the maximum pH which is the increase in potential V with each addition of O.1N KOH. . ++ +++ . . . Mn and Cr were not taken up by the sedlment Since there 18 no drop in pH below the pH of the pure sediment(see Table 1 & Figs 2,5) Fe++, C0++, Ni++, Pb++ show the greatest drops in pH. This means that these metals have the greatest tendency to be taken up by the + ++ . + . . sedlment. Fe + , Hg , and Ca++ also show a conSlderable dr0p in pH when they were added to the sediment, also implying a readiness to be taken up. Cu++, Zn++, and Cd++ Show the least drop in pH but unlike ++ +++ Mn and Cr , they did form complexes with the sedimenttTable 1 and Figs 1-6). In Fig.6, the titration curve of Hg++ shows two inflection points. Martell and Calvin(1952) provide an explanation for this(pg an). The first inflexion point represents the completion of the normal chela— tion of the metal: CHZCOO‘ + +2 --———4s M + HN<:::~CH2COO CHZCOO .__W -1 The second inflexion point indicates another reaction with base. This could be due to interaction with excess chelating agent to form a higher chelate: \"J “OOC CQ\\\\\\ ///// CH 0 / l —~-~ a ””1"\\“CH .’ /' 2 \\‘0 ———— 00 2 CHZCOO' + .__—.__) + HN ———— 0H2000‘ (.__ 0H2000 _.r4 -——— 000’ + H+ -12... Schwarzenbach(19u8), found that this behavior is characteristic of the divalent positive ions of Ca,Co,Cu,Fe,Mg,Mn,Ni,Hg,Pb and Zn as well. Relative Stability of the Metal Complexes: Mellor and Maley(19u7) have determined that the stability of complexes of bivalent metal ions follow the order: Pd >Cu>Ni >Pb >Co> Zn >Cd >Fe >Mn >Mg , irrespective of the nature of the ligands involved. The stability also seemed to decrease with increasing basicity of the metal. Irving and Williams(19u8), also corroborated the findings of Mellor and Maley(1947) on the successive stability constants of the divalent ions of the first transition series. This is generally: Cr>Mn (Fe-0o+f>>2n+f>.Ni++ which did not conform to the Irving-Williams series for stability constants of metals with ligands. This difference is probably due to the fact that humic acid is heterogenous in nature and has a compli— cated structure. Because of this, no exact mechanism of reaction between metal and humic acid could be ascertained. In my study, titration of humic acid was carried out but no inflec— tion points, from which the equivalent weight of humic acid could be calculated, were obtained. As Breger(1963) pointed out, the titration curve resembles the adsorption isotherms of colloidal substances. Later experiments however proved that this was true chelation rather than adsorption. Also in my study, titration of humic acid in the presence of 5,000mg each of Cu++, Co++,Zn++, Hg++ and 0d++ showed the trend of stability similar to the findings of Adhikari and Hadzra(1970): Cu++i>Co++>~Zn++. In addition, the metals were not precipitated during titration, indicating the formation of a soluble metal organic complex. Ovcharenko and G0rdienko(1969) also observed that the stability of the humic acid metal ion complexes deduced from the potentiometric curves for the titration (with 0.01N NaoH) of the acids in the presence of varied concentrations of metal salts was in the following + ++ . . . order: Fe++> Cu++> Co > Mn+ This too did not agree With the Irving—Williams series. Relative Stability Of Metal Complexes Of Lake Lansing Sediment Metal equISaIgnce giggiliiy point Hg++ 5.05 1.00 0o++ 5.20 1.03 Pb++ 5.35 1.06 Ni++ 5.40 1.09 Fe++ 5.85 1.16 Ca++ 6.40 1.27 Zn++ 7.55 1.50 Cu++ 8.65 1.71 Computation of Relative Stability: Relative order of stability of the metals which formed complexes with the organic sediment (divalent metals only) was calculated as follows: Hg++ has the least pH at the equaivalence point: therefore using the pH of Hg++ at equivalence point as a standard, we can divide the pH of the other metals at the equivalence point by the pH of Hg++ at equivalence point, to give the relative order of stability of these other metals. Hence the order of stability of 0o++ is 5.20/5.05 = 1.03. Relative stability deduced from pH at equivalence point in this manner shows the following order for the metals compared: Cu++> Zn++>Ca++> Fe++>Ni++>Pb++> Co++) Hg++ . This does not con- form with the Irving—Williams series in which Fe++ is the least stable. Although calculation of the actual stability constants of the metals using the Bjerrum method for example would be valuable,we cannot do this in this case because the molecular weights of the complexing agents in our sediment are not known. Hence the resort to relative stabilities which are also useful. _LJ- TABLE 3 Summary Of Relative Stabilities Of Humic Acid Fractions In Four Studies Order Of Stability Average molecular Wt- . of humic ac1d Adhikari & Hazra(1970) Adhikari & ++ ++ ++ .++ Hazra(1970) Cu :>CO Z>Zn 3>Nl 670 (2.494 2.453 2.126 1.385) Gordienko & Ovcharenko Fe++J>Cu+f>.Zn+I>.Mn++ ' (1969) Courpron ++ ++ _ Przifig; Cu++S>Co++C>Zn. - Displacement of One Metal By Another 5000mg of Fe++ plus SOOOmg of Cu++ were added to the sediment before titration. pH drop was 2.05. During titration with 0.1N KOH, Fe++ was precipitated at pH 6.05. Similarly, when Cu++ and Hg++ were added, the pH dr0pped to 3.45. Explanation The metal that forms the most stable species complexed first, but Fe++ and Hg++ have a greater tendency to form chelates with the sedi- . . . . + ment: hence the conSlderable dr0p in pH. On titration, Fe++ and Hg + were precipitated showing that Fe++ and Hg++ have been displaced from the sediment into the overlying water. Summary and Conclusion The results of this study show that Fe++, C0++, Ni++, and Pb++ have the greatest tendency to form complexes with the organic sediment used. The relative stabilities are of the order: (:t.1++>zn‘“i> Ca++> Fe++> Ni++ >131)“; Co”) as”; 0d”) Hg“; Fe+++> Fe++ In defining relative stabilities, the important factor is the pH at equivalence point. We assume that all the metal complexes, in the same environment, are initially at the same pH level. Thus as pH increases, the metals are differentially pushed out of the metal com- plexes are either precipitated as insoluble hydroxide or remain in solution. The eXperimental procedures used in this study could be used in defining the profile of a lake: specifically to find out the capacity for metals of a Lake sediment. It is thus applicable to environmental pollution studies, and in explaining some aSpects of ore formation, and why one metal is not found in association with another. APPENDIX A Test For Reproduceability Of The Method Used The pure sediment was titrated with 0.1N KOH, three times. The 4% 95% confidence level for each of the points is repoted below: TABLE A Confidence Level Of Titration Points pH Readings Confidence Level Vol.0f 0.1N —m”'1st ’2nd 3rd KOH added Titration Titration Titration EER X ACER) 0.00 6.85 6.85 6.95 0.10 6.883 I 0.1301 0.10 8.35 8.35 8.35 0.00 8.350 I 0.0000 0.25 9.65 9.75 9.65 0.10 9.683 I 0.1301 0.50 10.15 10.15 10.15 0.00 10.150 I 0.0000 0.75 10.35 10.30 10.45 0.15 10.367 I 0.1952 1.00 10.60 10.u5 10.55 0.15 10.533 I 0.1952 1.25 10.75 10.55 10.75 0.20 10.683 I 0.2602 1.50 10.85 10.80 10.85 0.05 10.833 I 0.0651 2.00 11.00 10.95 11.00 0.05 10.983 I 0.0651 3.00 11 20 11.15 11.20 0.05 11.183 I 0.0651 n.00 11.30 11.30 11.30 0.00 11.300 I 0.0000 5.00 11.u0 11.n0 11.u5 0.05 11.417 I 0.0651 * The formula used for calculating the 95% confidence level (three observations per point) was: CL = x I A(ZR) where X is the mean of the three observations, A is the confidence interval constant for three observa- -tions per point, and, 57R represents the range of the observations per point. APPENDIX B Titrations Of Sediment Sample (From lst Collection) With O.1N KOH. In The Absence Of Metal Ions. 2. In the presence of metal catiopg. Vol.0f 0.1N KOH pH Vol.0f 0.1N KOH pH added (in mls) Readings added (in mls) Readings 0.0 7.55 With 5000mLof Cu++ 0.1 8.35 ' 0.25 9.55 0.0 6 35 0.50 10.15 0.25 6.55 0.75 10.45 0.50 6.75 1.00 10.60 0.75 7.00 1.25 10.75 1.00 8.65 1.50 10.85 1.50 10.35 2.00 11.05 2.00 10.75 3.00 11.20 3.00 11.15 4.00 11.35 4.00 11.25 5.00 11.45 . . . +++ 2nd tltratlon Wlth SOOCyg Cr 04) 7.60 0.0 8.05 0.1 8.45 0.1 8.80 0.25 9.75 0.25 9.85 0.50 10.15 0.50 10.35 0.75 10.45 1.00 10.85 1.00 10.60 1.50 11.05 1.25 10.75 2.00 11.15 1.50 10.85 3.00 11.40 2.00 11.00 4.00 11.40 3.00 11.20 ++ 4.00 11.35 With 5000pg_Cu and r 3rd titration 0.0 6.55 0.25 6.90 0.0 7.55 0.50 7.40 0.1 8.30 1.00 10.05 0.25 9.80 1.50 10.85 0.50 10.20 2.00 10.90 1.00 10.65 3.00 11.25 1.50 10.85 4.00 11.35 2.00 11.05 3.00 11.25 4.00 11.35 5.00 11.45 ii Vol.0f O.1N HOH pH Vol.0f 0.1N KOH pH added (in mls) Readings added (in mls) Readings With 5000pg Mn++ 21.00 3.65 ‘ 22.00 4.50 0.0 8.05 22.50 5.20 0.25 9.30 23.00 6.00 0.50 10.05 24.00 7.75 1.00 10.35 25.00 9.85 1.50 10.65 26.00 10.50 2.00 10 75 27.00 10.75 3.00 10 90 28.00 11.00 4.00 10 95 29.00 11.05 5.00 11.15 30.00 11.15 With 5000pg Ni++ With 5009pg Pb++ 0.0 2.25 0.0 2.45 0.25 2.25 0.25 2.45 0.50 2.25 1.00 2.45 5.00 2.25 5.00 2.65 10.00 2.25 6.00 2.85 12.00 2.35 7.00 3.10 13.00 2.35 8.00 3.50 15.00 2.40 9.00 5.35 17.50 2.45 9.50 6.35 19.00 2.55 ' 10.00 8.80 20.00 2.60 10.50 9.75 21.00 2.65 11.00 10.25 22 00 2.75 12.00 10.65 23.00 2.85 13.00 10.95 24 00 3.00 14.00 11.05 25 00 3.15 15.00 11.20 26 00 3.30 20.00 11.45 S; 2% 2:28 With 5000pg of Fe+++ 28 00 5.40 28.50 6.25 8'35 5'3; 29.00 6.55 0°50 3°65 30.00 8.00 0'75 4'65 31.00 9.30 1°00 5°80 32.00 10.00 1°25 7°05 33.00 10.80 1'50 8°50 34.00 11.05 1'75 9'20 35.00 11 25 2:00 9:65 With 5OOCpgACo++ 3.00 10.50 ' 4.00 10.80 0.0 2.15 5.00 10.90 0.25 2.15 6.00 11.05 0.50 2.15 1.00 2.15 2.00 2.15 5.00 2.25 10.00 2.25 15.00 2.50 20.00 3.15 (contd. on right) Vol.0f 0.1NKOH pH Vol.0f 0.1NKOH pH added (in mls) Readings added (in mls) Readings W' II W'th 500C 5 0a++ lth 500Cpg Fe 1 P. 0.0 2.05 0.0 3.50 0.25 2.05 0.25 3.65 0.50 2.05 0.50 4.10 1.00 2.05 0.75 4.95 2.00 2.05 1.00 6.40 3.00 2.05 1.25 8.05 5.00 2.10 1.50 9.55 8.00 2.15 2.00 10.20 9.00 2.20 3.00 10.75 10.00 2.25 4.00 11.05 13.00 2.40 5.00 11.20 .00 2.55 . 3 ++ i8.00 2.95 with 500Cp3 of Cd 19.00 3.15 0.0 7.05 20.00 3.65 0.25 8.50 21.00 5-85 0.50 9.45 22.00 7.95 1.00 10.30 23.00 9.80 1.50 10.55 24.00 10.45 2.00 10.90 25.00 10.70 3.00 11.05 26.00 10.90 4.00 11.20 27.00 11.00 5.00 11.25 30.00 11°25 With 5000pg Hg++ With 5OOCpg Zn++ 0.0 4.30 0.0 6.35 8'é5 g°gg 0.25 6.65 0°50 7°20 0.50 7.15 0'75 7°65 0.75 7.55 1'00 7'95 1.00 9.20 1°25 8°85 1.25 9.90 1°50 9°80 1.50 10.20 2°00 10'50 2.00 10.50 3'00 10°85 3.00 10.85 4°00 11'05 5.00 11.15 ° ' 5.00 11.15 . ++ . ++ With 5000gg Fe and With SOOOpg Hg and 5000pg Cu SOOCPE Cu++ 0.0 2.10 0.0 3.45 5.00 2.15 0.25 3.95 10.00 2.25 0.50 4.45 15.00 2.60 0.75 5.20 16.00 2.75 1.00 6.15 18.00 2.95 1.50 6.85 19.00 3.10 1.75 7.70 20.00 3.40 2.00 8.85 21 00 4.60 3.00 10.60 22 00 6.05 4.00 10.90 23.00 8.20 5.00 11.05 24.00 10 65 ' 25 00 10.70 30.00 11 20 (I, Titration Of Sediment, 2nd Collection. (for testing reproduceability of method ) Vol. of 0.1N KOH added (in mls) pH Readings In The Absence of Metal Ions 0.0 .10 .25 .50 .75 .00 .25 .50 .00 .00 .00 .00 \ni—‘WNHHHOOOO 2nd titration .0 .10 .25 .50 .75 .00 .25 .50 .00 .00 .00 .00 \n-P'KJJNHHHQQQQQ 3rd titration .0 .10 .25 .50 .75 .00 .25 .50 .00 .00 .00 .00 \R-P'WNHHHOOOOO 6. 8. 9. 10. 10. 10. 10. 10. 10. 11 11. 11 85 35 75 15 30 “5 55 80 95 .15 3O .40 .85 .35 .65 .15 .35 60 .75 .85 .00 .20 .30 .40 Titration 0f Humic Acid. (10gm Humic acid in 1000mls H20: 50mls aliquot used for each titration) Titration In Absence Of Metal Ions. Vol.0f 0.1N KOH pH Vol.0f 0.1N KOH pH added (in mls) Readings added (in mls) Readings 0.0 3.65 8.00 8.25 0.10 3.75 9.00 8.65 0.25 3.85 10.00 9.45 0.50 4.00 11.00 10.00 0.75 4.15 12.00 10.45 1.00 4.25 13.00 10.95 1.25 4.45 14.00 11.05 1.50 4.75 15.00 11.25 1.75 4.90 gzgg 2:?2 With 5000pg Co++ 2.50 5.35 0.0 2.35 3.00 5.75 0.50 2.35 3.50 5.90 2.00 2.35 4.00 6.25 5.00 2.35 4.50 6.55 7.00 2.40 5.00 6.70 starts dissol- 10.00 2.45 5.50 6.90 ving.pH reading 13.00 2.55 6.00 7.05 unsteady.Takes 15.00 2.65 6.50 7.15 up to 5 mins. to 17.50 2.65 7.00 7.25 get constant 20.00 2.85 8.00 7.75 reading. 22.50 3.05 8.50 7.95 25.00 3.45 10.00 8.75 27.00 4.15..diss0— 15.00 11.25 28.00 4.75 lution ++ 29.00 5.35 starts With 5000pg Cu 30.00 5.95 31.00 6.55 0,0 3.45 32.00 7.45 8'2 3.48 33.00 8.45 0:56 5.55 34.00 9.10 1.00 3.75 35°00 9'40 1.50 “.05 40.00 11.25 2.00 4.25 2.50 5.25 2.75 5.45 3.00 5-75 3.25 6.05 3.50 6.15 3,75 6.25 4.00 6.55 4.50 6.85 5.00 6.85 5.50 7.25 dissolution 6,00 7.90 starts 6.50 7.70 7.00 7.85 (finn+A A“ “:~k+\ Titration with 5000mg Zn++ Vol.0f 0.1NKOH added (in mls) .O .25 .50 .00 .OO .00 .50 .00 .50 .OO .50 .OO .00 .00 .00 .OO .00 .00 .OO .00 raner4 \J’IP’HOOCDVO‘xU‘U‘ktl-‘KJKONHOOO N O with SOOCpg Hg+ 0.0 .50 .OO .00 .00 .00 .OO .00 .OO .00 .00 .OO .00 .OO .00 .OO .00 .00 H O\OCI)\}O\U\4:'\JONHO NrehnareH OCDUxPTUP‘ + pH Readings iareH HCDCnombarq\aomhownvun$>¢\ounbxo Hear- HHO\O\O CDCDV‘V} O\O\\J‘\U'\ {-‘wwww .55 .55 .65 .75 .35 .95 .35 .65 .95 .20 .50 .75 .45 .80 .65 .00 .15 .50 .65 .35 solution starts solution starts With 5000mg of 0d++ Vol.0f 0.1NKOH added (in mls) Hra U10 \OCDVQU‘PWNHOO .O .50 .00 .OO .00 .00 .OO .00 .OO .00 .00 .00 .OO P‘H pH HO \O\Om\lO\U\U\ F'LQKAJLO Readings solution starts to form BIBLIOGRAPHY Abelson,P.H.: Researches In Geochemistry. 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Agrochem Talajpan, (Hungary), 18(1), 25-33. 1969 Ringbom,A.J. Complexation in analytical chemistry: a guide for the selection of analytical methods based on complexation reactions. New York, Interscience Publ. 1963. Schnebetkovskii & Yu. V. Kuznetzsov : " Behavior of Cs-137 and Ce-144 in the sorption system of sea-water sediment" Translated from Radiokhimiya, vol. 13, no.6, 911-913, N0v-Dec.1971. Stanton,R.L. Ore Petrology. McGraw-Hill. 1972. Titley,S.R. Hicks,C.L. Geology_of the porphyry copper deposits, South Western North America. Univ.of Arizona Press, Tucson, Arizona, 1966 CH IGAN STATE UNIVERSITY LIBRARIE IIIIIIIIO III III I III I