THE NATURE OF CALCIUM - ORGANIC COMPLEXES IN NATURAL SOLUTIONS AND ITS IMPLICATIONS ON MINERAL EQUILIBRIA Thesis for the Degree of Ph. D. MICHIGAN STATE UNIVERSITY JAMES I. HOFFMAN 1969 _”— —-w L IB RA R Y Michigan State University This is to certify that the thesis entitled The Nature of Calcium-Organic Complexes in Natural Solutions and its Implications on Mineral Equilibria presented by James I. Hoffman has been accepted towards fulfillment of the requirements for Ph.D. Geology degree in QM/MJ Major professor Date July 18, 1969 0-169 ABSTRACT THE NATURE OF CALCIUM - ORGANIC COMPLEXES IN NATURAL SOLUTIONS AND ITS IMPLICATIONS ON MINERAL EQUILIBRIA BY James I. Hoffman Dissolved organic complexes are demonstrated to play such an important role in many natural aqueous environments that present, purely inorganic models must be severely modified. In many environ- ments, especially in sediment pore waters, organic complexes proba— bly act as the controlling factor for solution or precipitation of min- erals. Organic complexes cause cation activities to behave as compli- cated functions of Eh, pH, and the mass budget of dissolved organic matter. Presence of these complexing agents can cause solution of minerals whose cation‘s concentration is at a level of saturation. Natural samples, containing the entire spectrum of available organic and inorganic components, were investigated to determine the role of dissolved organic matter in mineral equilibria. The magnitude and the effects of organic complexes on calcium ion activities were evaluated for surface waters and sediment pore waters from central Michigan lakes and streams, and marine and brackish waters of coastal North Carolina. Strong organic complexes were found to occur in sig- nificant proportions in many of these environments with as much as 50% of the total calcium existing in a complexed state. James I. Hoffman The nature of organic complexes was investigated by monitoring calcium activities with a calcium activity electrode and comparing results with the total concentration obtained by atomic absorption spec- trophotometry and observations of pH, while various artificial pertur- bations were performed on natural fresh water samples. Within samples, relationships between micro-organisms,atmos- pheric oxidation, and dissolved organic matter caused increases in pH and calcium activity with time. Thus, natural samples must be ana- lyzed immediately or valuable information will be lost. Additions of excess amounts of smaller divalent and trivalent cations caused ion exchange reactions between the calcium-organic complexes and added cations. Exchange preferences were not entirely in accordance with predictions based on present theories. The role of organic complexes in weathering takes on new implications, because slightly soluble cations such as iron and copper can now be effectively solubilized and transported as organic complexes. Artificial additions of acid caused release of bound calcium ions. Interactions of added hydronium ions with dissolved organic matter produced strong pH buffering reactions between pH 7. 6 and 6. 6. These processes were irreversible. Slight increases in the oxidation potential created new com- plexes which lowered calcium activity; but strong oxidation destroyed Organic matter and released bound cations. A new aqueous geochemical model, needed to characterize the James I. Hoffman effects of dissolved organic material, was constructed from observa- tions Of a sample environment. The descriptions of ion activities were quasi-linear; those for pH variations were reciprocally linear. Future, universal generally acceptable models will have to be empirically de- termined from natural samples, whereas the current inorganic models are derived from synthetic systems. Gathering empirical data must involve determination of rates of influx and concentrations of natural acids, organic acids, dissolved organic matter, oxidizing and reducing agents, and the roles of geochemically active micro-organisms. Col- lection of these data will pose a formidable task, and interdisciplinary efforts will be needed for precise development of this ultimate model. THE NATURE OF CALCIUM - ORGANIC COMPLEXES IN NATURAL SOLUTIONS AND ITS IMPLICATIONS ON MINERAL EQUILIBRIA By (7 I Jame s If Hoffman A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Geology 1969 ACKNOWLEDGMENTS Extraordinary acknowledgment must be given to Dr. Robert Ehrlich. More than anyone else his advice and stimulation have shaped my concepts of scientific thought into their present form. His critical evaluation of this investigation, confidence in the re- searcher, and tireless efforts were invaluable in guiding this project to successful completion. Special thanks is given to Dr. Samuel Romberger, who was responsible for my training as a geochemist. His consultations and suggestions during the course Of the investigation, and his many editorial efforts on this manuscript were most helpful. Many other individuals have offered encouragement and help. Dr. Thomas Vogel and Dr. Hugh Bennett have both added some mean- ingful measure to the completion of this investigation. Not to be over- looked was the unceasing support of my wife Ellen, and her considerable aid in the final preparation of this manuscript. ii TABLE OF CONTENTS List of Tables ---------------------------------------------- v List of Figures --------------------------------------------- vi Introduction ------------------------------------------------ 1 Experimental Design ---------------------------------------- 7 Introduction -------------------------------------------- 7 Preliminary Experiment --------------------------------- 8 Instrumental Design ------------------------------------- 14 Instrumental Methods --------------------------------------- 22 Results .................................................... 34 Introduction ............................................ 34 Relative and Absolute Magnitudes of Organic Complexification - 34 Fresh Water Environments -------------------------- 34 Marine and Brackish Water Environments ------------- 39 Time Degradation --------------------------------------- 42 Perturbation Experiments ------------------------------- 45 Introduction ---------------------------------------- 45 Cation Exchange ------------------------------------ 47 pH Perturbations ----------------------------------- 50 Oxidation Perturbations ----------------------------- 57 iii 3%...- Table of Contents (cont'd.) Role of Organic Complexes in Calcite Equilibria in Burke Lake - 64 New Aqueous Geochemical Model ----------------------------- 70 Introduction ............................................ 70 Example of the New Model ------------------------------- 72 Summary and Conclusions ----------------------------------- 77 List of References ------------------------------------------ 80 References Cited --------------------------------------- 80 General References ------------------------------------- 81 Appendix A Preliminary Experiment ------------------------- 87 Appendix B Calcium Ion Activity in Parts per Million ---------- 88 Appendix C Ion Exchange Experiments ----------------------- 89 Appendix D pH Perturbations ------------------------------- 92 Appendix E Eh - pH Perturbation Experiment ----------------- 102 IV LIST OF TAB LES Major Ionic Species in Michigan Sampling Sites ------------- 13 Instrumental Analysis of Variance ------------------------- 18 Time Degradation Experiment ---------------------------- 31 Magnitude Of Calcium Complexification in Fresh Waters ----- 35 Magnitude of Calcium Complexification in Marine and Brackish Waters ------- 40 Time Degradation of Red Cedar River Sample -------------- 45 Eh - pH Interaction Design ------------------------------- 60 Calcite Equilibria Data for Burke Lake -------------------- 66 10. ll. 12. l3. 14. 15. LIST OF FIGURES Central Michigan Sampling Sites -------------------------- 10 Coastal North Carolina Sampling Sites --------------------- 15 Instrumental Variation Trends ---------------------------- 19 Infrared Spectra of Red Cedar River Residue --------------- 24 Time Degradation Effects on Complexed Calcium Ions --------- 43 Release of Complexed Calcium vs. pH with Time ----------- 46 Ion Exchange Capacity of Organic Complexes --------------- 48 Acid Titration of Fresh Water Lakes ---------------------- 51 Acid - Base Titration of Red Cedar River ------------------ 53 Calcium Activity - pH Relationships ----------------------- 56 Calcium Activity - Eh Relationships ----------------------- 59 Eh - pH Effects on Calcium Activity ----------------------- 62 Eh - Acid Effects on Observed pH ------------------------- 63 Calcite Equilibria in Burke Lake -------------------------- 68 Observed Calcium Activities vs. New Model ---------------- 74 vi 1 _ <-&§§ ._ ,, ... - INTRODUCTION The earth‘s surface has been undergoing evolutionary changes for billions of years. Much of the change has been geochemical in na- ture, resulting from interactions between the low temperature aqueous surficial environment and materials formed under different equilibrium conditions (many characterized by high temperatures and pressures) and brought to the surface by various modes of crustal dynamics. Also, owing to the dynamics of the earth as a planet and its relative position in the solar system, environmental conditions on the continental sur— faces do not remain constant for great lengths of time. Thus, materials that reach chemical equilibrium under surficial conditions at a given time, may be out of equilibrium at a later time. Water serves three important functions in chemical processes in a natural environment: (1) it is a chemically reactive substance in its own right; (2) it serves to dilute or concentrate other reactants and products; and (3) it is a medium for other reactants and products. One of the major goals of low temperature aqueous geochemis- try is to determine the chemical reaction sets which constitute the sys- tems of weathering, solution, and deposition in order to delineate their critical controlling parameters. More specifically, aqueous geochem- istry aims to define ion stabilities and mineral equilibria to explain and predict their behavior in the natural environments. Chemists have long known that certain parameters tend to mod- erate aqueous reactions. Changes in the hydronium ion content (pH) resulting in acid-base type reactions, and trends resulting in the gain or loss of electrons which dictate the oxidation potential (Eh) are very important. A third variable, ionic strength, influences individual ion activity* which directly affects the quantity of a particular ionic species in solution. These three factors form the basis for current geochemi- cal models. Current research in low temperature aqueous geochemistry has been primarily concerned with the behavior of simple inorganic ana- logues of natural systems. That is, the nature of the interactions of a relatively small number of ions under various conditions of Eh and pH have been studied in synthetic solutions. This simplified synthetic ap- proach, summarized by Garrels and Christ, (1965), has been deemed necessary due to the inherent complexity Of the natural system. Results Of this approach are not adequate to entirely explain ei- ther the nature or magnitude of observed weathering, transport, and precipitation phenomena or those inferred from the nature of products in the geologic record. For instance, conditions needed to explain Precambrian iron formations require the simultaneous transport and >I‘activity — is the thermodynamic concentration, which is usually less than the total concentration in surface environments. It is the portion of the total concentration which is free to participate in chemical reactions. precipitation of large amounts of iron and silica. Since, in terms Of standard mineral equilibria theory, iron is most mobile under pH con- ditions that are unsuitable for concomitant transport of silica, current theory results in either a seeming paradox or an acceptance that cur- rent geochemical models are seriously incomplete. The failures of the current model may be due to incomplete evaluation of all possible factors influencing ion solubility and trans- port. A factor which has not been much evaluated is the role dissolved organic substances play in aqueous environments. Because living or- ganisms, both plant and animal, need elements in ionic form for sur- vival, living matter has effectively increased weathering and solution reaction rates many fold over similar inorganic reactions. In addition to solubilization in response to nutrient needs, the earth‘s surface is bathed in organic metabolic wastes and the degradation products of dead organisms. It is surprising in this light that organic-inorganic interactions have not been more widely studied. A few investigators over the past fifty years have indicated the relevance of dissolved organic substances in geochemical models. Important among them is the work of Gruner, (1922), who recognized the transport problems inherent in the nature of siliceous iron forma- tions. He demonstrated that relatively large amounts of iron and silica could be transported in organic rich bog waters. Strengthening his case, Birge and Juday, (1934), showed that large amounts of dissolved organic matter were commonly present in fresh waters. 4 In general, current research concerning soluble organic matter has been directed along three approaches: (1) qualitative observations of natural systems; (2) observations of i_n_v_it_r_c_> reactions between rnin- erals and known organic species present in natural waters; and (3) pre- cise identification of the dissolved organic species present in natural waters. Baas Becking, et a1, (1960) corroborated earlier more qualita— tive works in terms of Eh and pH limitations of natural environments, and showed that organic substances are present over a much broader range of Eh and pH than was previously expected. This type of work on general characteristics of natural waters is cruder than _i_n_li_’£r_o laboratory experiments, but has direct relevance to understanding the workings of natural aqueous systems. Other researchers have demonstrated the effectiveness Of reac- tions between soluble organic matter and common minerals. Evans, (1964), showed that quartz, carbonates, clay minerals, and ferromag- nesian silicates could be rapidly solubilized by ATP, an organic com- plexing agent found in cell tissue. Bailey, (1967), showed that oxalic acid would effectively weather ferromagnesian minerals in a short time. Hem, (1965), showed that manganese was complexed and trans- ported by gallic acids. Shapiro, (1964), and Christman, (1967), have shown that yellow organic acids can complex and transport a variety of cations. These studies have shown the importance of specific organic substances in specific reactions, but the reagent strengths used experimentally and the simplicity of their systems are geologically un- reasonable. All of these researchers designed their experiments to avoid the complexity and potential interactions which occur in natural environments; therefore, little information was gained concerning the relationships between Eh, pH, and dissolved organic substances. Many workers have documented the presence of particular dis- solved organic species in natural waters. Several hundred organic compounds have been identified; common among them are the dicar- boxylic acids (such as oxalic and citric acid) which are effective com- plexing agents. Many of these soluble organic species and the experi- mental techniques used in identification are summarized by Breger, (1963), and Colombo and Hobson, (1964). Such data are helpful, but two factors combine to limit their usefulness: (1) species are identi- fied without relative percentages and absolute concentrations; and (2) the analytical procedures so greatly modify the natural sample that in many cases the compound identified probably is a degradation product created by the separation procedures used prior to identification. This study, modeled in spirit after Gruner's, (1922), is an at- tempt to understand the role of dissolved organic substances in mineral equilibria by observations of the nature of cation-soluble organic reac- tions in natural samples. The great variety Of dissolved organic mat- ter, with its wide spectrum of chemical properties, presents an enor- mous number of potential interactions with the inorganic portion of the geochemical environment. This study is designed to illuminate the nature of the sum effect of the organic species, and their interactions on solution equilibria. Thus, for purposes of this study, the dissolved organic system has been observed and manipulated. Several aspects of the role of organic complexes in mineral equilibria must be evaluated: (1) the quantitative importance of soluble organic complexes in aqueous environments; (2) the stability of these complexes; and (3) what effect interactions between soluble organic mat- ter, Eh, pH, and cation composition has on the solution, transport, and deposition of minerals. To solve these problems, this study has used natural samples in order to evaluate the interactions between dissolved organics and inorganic components and parameters. Analysis has been made on complete natural samples to characterize the role of organic complexes in the aqueous environment, rather than relying on systems containing only a few components in a less than real analogue. It is the combined behavior of the entire aqueous melange which is geologically significant. EXPERIMENTAL DESIGN Introduc ti on The purpose of an experimental design is to plan the logic of an investigation to insure sufficient investigational coverage for solution of the problem. This presumes a subsequently derived clear formula- tion of the problem. The design insures, in an experimental sense, that the previously determined scope and relevance of the problem not be compromised. It provides an estimate of the magnitude of the task needed to solve the problem; and allows the investigator to plan for op- timal efficiency of the research with regard to time and resources. The first step in this study was to define the problem and eval- uate its relevance to major geologic problems. This was done by con- sideration of the general role of organic substances; evaluation of the ideas of past workers; and by conducting a preliminary experiment to determine if, in fact, organic complexes of common cations existed in significant enough quantities to be observed with adequate precision with the means at hand. The first two points were evaluated and dis- cussed in the Introduction. Preliminary Experiment The preliminary experiment was undertaken to discover whether the effect was of such a nature that it could be evaluated with our ana— lytical resources. It was important that the experiment yield some qualitative measure of the magnitude of the complexing effect, in order to estimate its importance in aqueous geochemical environments. The experimental technique was designed to investigate the presence of nat- ural organic complexes by placing them in competition for calcium ions with a synthetic complexing agent, EDTA. This technique involves the removal of divalent calcium and magnesium ions from solution by binding them to the EDTA molecule. This process, called chelation (Martell and Calvin, 1952, pp. 9-15) results when complexes are formed between the polydentate ligand and the cation. EDTA is an extremely effective synthetic chelating agent for divalent cations. Eriochrome Black T, the colored indicator, is a weaker chelating agent, which displays a red color when it sequestors calcium and magnesium ions, and a blue color in the uncomplexed state. In the analysis, as EDTA is added during titrations, it first extracts all free calcium and magnesium from solution. Next it depletes those cations weakly bound to possible natural organic com- plexes; and finally the EDTA robs the indicator of its cations. In response, the Eriochrome Black T changes to a blue color. If com- plexing or chelating agents stronger than the indicator are present, 9 the EDTA will not have depleted them of their bound ions when the indi- cator turns blue. Thus some complexed calcium and magnesium in solution would escape detection. Since inorganic solutions contain no complexes which can compete with Eriochrome Black T, any such bound portion escaping detection must be in combination with dissolved organic matter. This conclusion is strengthened if that bound fraction can be released by destruction of the organic complexing agents. A sample was taken from the Red Cedar River, (site B, Figure l) and split into six 50 Inilliliter (rnl.) aliquots which were analyzed by the colorimetric EDTA - Eriochrome Black T technique for total calcium and magnesium (Martell and Calvin, 1952, pp. 482 - 490). Three of six aliquots of sample were analyzed by this method for total calcium and magnesium. The other three aliquots were first evaporated to dryness; 10 ml. of concentrated nitric acid was then added to each residue to oxidize and degrade any organic matter present; and the resultant was slowly evaporated to dryness. After the second drying, the dried residue was dissolved in dilute hydro- chloric acid and the solution was brought to its original aliquot volume with distilled water. These aliquots were then analyzed for total cal- cium and magnesium. Results (Appendix A) indicated that the total calcium and mag- nesium present in the untreated samples was about 112 parts per mil- lion (ppm) whereas in the oxidized samples, the total for these ions was Figure / CENTRAL MICHIGAN SAMPLING SITES I """" | Lake Lansing 11 about 165 ppm. Since no calcium or magnesium was added in the course of the analysis, one must conclude that the presence Of a sig- nificant proportion (about 30% in this case) of these cations was chem- ically masked in the untreated samples. Further it must be concluded that the aforementioned cations are strongly bound to an oxidizable host, because no known inorganic model can explain such results. These results, along with previous preliminary concepts help define the nature of the ensuing investigation. Factors such as the interactions between dissolved organic complexing agents, oxidizing agents, and hydronium ions are of primary interest. These interac- tions were investigated by monitoring cation activities with changes in the Eh and pH of the samples. Secondly, variations of magnitudes Of organic complexification of the common cations were investigated in both free waters and pore waters in fresh water and marine environ- ments to determine the ubiquity of the effect. Finally, the current model for aqueous geochemical systems has been modified to take into account the role dissolved organic mat- ter plays in mineral equilibria in these environments. This involved expressing conclusions in a manner suitable for use in geochemical calculations. A variety of geochemical environments are necessary to inves- tigate this problem. Fresh waters in central Michigan (Figure 1) were chosen as the primary sampling sites and research efforts were con- centrated here, because the ionic strength is low enough that no 12 inorganic complexing occurs to confuse the data. Since these samples may be treated as thermodynamically ideal solutions, any variations in cation activities in undersaturated solutions with changes in the im- portant geochemical parameters must have been due to their interac- tions with dissolved organic complexes. A reconnaissance study was made of several Mid-Michigan lakes in order to choose sampling sites which would exhibit a diversity of chemical compositions and environments. Data for concentrations of geologically important ionic species were collected at the sampling site using a Hach Chemical Company Direct Reading Engineer‘s Laboratory in accordance with their standard experimental techniques (Table 1). After reviewing these data, Holt Lake was excluded as a sampling site because of its chemical similarity to the Red Cedar River. The Red Cedar River became the primary sampling site be- cause it was known to have artificially enhanced organic levels (pollu- tion), and because the preliminary experiment indicated that there was a significant interaction between dissolved organic matter and calcium and magnesium ions. Several shallow lakes proximal to the laboratory were chosen representing a range of fresh water environments. Burke Lake is a Small ground water fed system with little inflow of surface water and no man-made organic pollutants. Three environments within this lake were investigated: (1) pore waters from the upper five centimeters of 13 00 .oom 0m .0 0% .m 00 .mm 00 .ow 0N. .h 334 Omom .30H .m .mo .9 .mooi £3.30 was .39”an Eon.“ woumfiomuofi: one? mucoaodmooo L» .mumv of Eon.“ 68.35030 one? mzuwnouu .GBOLm noflmhunoofloo mo oxsoa was exam Cook/awn Ohm mOSGHm .moumofimon mouse mo some OH: «nomenmou 00.Nm 00.NN 00.0 00 .0: w0.0 003v $0.0 00¢; 00.mm owlm wnwmcmd Oxmd NI. mod 00.0 mod OHNNA 00.0N 00.wN 00.0H 00 .mNN m0.0 00.mH 00.0 004m 00.wm 0w .N. 334 335mm 00.0 00 .0 00 .0 mica on m4 00.0wV 00.0NH 00.0H 00.00N 00.0 mm .0 00.m~ comm 0H .w Oxmd 303 No.0 00.0 No.0 NIOH N 3»; 00 .mm 00 .omH 00 JEN 00 .mm 00 .00a owfi uokfim a £50 Om m 030555 £0238 you 3an E mcofimficooaou H 3Com. mofim wQSQENm Gmwfioflz Hmnucoo aw mowoomm OAGOH uoflmz 00.0 00.0 00 .0 ~-ozs.ws.a 006v 00.0mH 00.0H 00.mm~ 00.0 00.: mod 00.mm 00 .m0H Hm .w AOL/Mm HmpOO pom :39ko m UMGOH «noonb mogm> mOO mour mo Sumcouum OMGOH HO eOm mOO mOom so am an m2 aO mm l4 sediments; (2) free waters proximal to the sediment - water interface; and (3) free surface lake waters. Surface waters of two other lakes were investigated because they showed degrees of eutrophication different from Burke Lake. Rose Lake is in a more advanced state of eutroPhication than Burke Lake, with only small amounts of artificial organic enhancement. Lake Lansing is a system with higher levels of both naturally and artificially enhanced dissolved organic matter than is found in Burke Lake; but lower levels than were found in the Red Cedar River during this experiment. To test the generality of the fresh water observations, comple- mentary, but more limited investigations of marine and brackish water environments were carried out in the vicinity of Topsail Island, North Carolina (Figure 2). Free water and sediment pore water samples were taken over a five day period in April, 1969, from both intra- barrier island marsh and from lagoonal environments. Instrumental De Sign Analytical techniques used in the preliminary experiment yielded only relative proportions of bound or complexed cations to free cations. To gain further information, activities of individual cations were monitored and related to changes in the important geochemical parameters. 15 3:6 q -1 _I Z 7/ 94.853 3&2: .\ 35 goat >82 mmtm oz_._a_2 nOHoanbmHQ AH come/H HO mOOthQ .HO OOhsom oHan mocmHnmxr HM mHmNHmn< 72 Na: 0.3 Na: Na: ~.$ son 33. 32.3 0 .4w w .8 e .8 o .vw o .ew o .mw swam soreness/w canons. m .mH m NH m .wH m .wH m .wH m .wH BOA £5.30on noH 0 .mm 0 .mm m .ow m .ow s .sw m .ow swam 3:83 8328 0 m Hy m N H mnofimauno Ono 0 Human OH OquHHmom GOH EDHOHmO unmadupmnH SRO moamHnw> .Ho mHmLHHdew HancoesbmzH N MOHAHsH l9 mOGOHH. :OHHmemxr HmunogdtmGH "m onstnH Ema nH ASHbHHOHw GOH EsHonO 4. O. O. 'l 00 .I .00 ‘l 0v L J1 0N H+ N+ uopenueouog uoI nunroteg umoux Luoxj mdd u; notietx'eA 20 known to exist in the Michigan sampling sites. The standard deviation for A. A. was 0. 69 ppm calcium, and for the calcium activity electrode this was 1. 66 ppm. These values were then used to determine the requisite sample size. Determination of the sample size needed to insure the sta- tistical validity of the results is perhaps the most important part of experimental design. The number of replications per sample is re- lated to the size of real difference one wishes to be able to detect as significant. It was decided that if 10% of the total calcium in a sample was organically bound, then this would represent a geologically sig- nificant portion, and variations of about 10% were of interest. This amounts to from 3 - 10 ppm in the fresh water samples (Table 1). Owen, (1967, p. 44), gives the relationship between requi- site sample replication needed to see a critical difference in a normal distribution with the probability of rejecting a true hypothesis (Type I error) of O. 05 and specifying the probability of accepting a false hypoth- esis (Type II error). Using the standard deviations determined from data gener- ated for the factorial analysis of variance, and choosing a critical difference of 5 ppm, a sample size involving three replications assured a Type II error of less than O. 01 for A. A. , and between 0. 05 and O. 20 for the calcium specific ion electrode. Review of Table 2 indicates that the average difference seen by each method is over 15 ppm cal- cium for the Red Cedar River. With a number this large, Owen's 21 charts indicate the chances of committing a Type 11 error with the calcium electrode fall to less than 0. 01. In summary, the experimental design has isolated the prob- lem and planned the attack. The preliminary experiment indicated the importance of organic complexes in fresh water. The sampling plan was formulated to test the ubiquity of organic complexes in a variety of environments and to investigate their relationships to important aqueous geochemical parameters. The instrumental methods of ana- lysis needed to solve the problem were compared and contrasted by a factorial analysis of variance. Standard deviations were calculated for instrumental precision and for determination of the requisite sample size used in data collection. INSTRUMENTAL METHODS Having formulated the experimental design, the following standard routine for sample analysis was established. The samples were: (I) collected in polyethylene containers; (2) filtered for all particles greater than five microns in diameter; and (3) split into aliquots for analysis by A. A. , specific ion electrodes, and pH. All water samples were collected in 500 ml. or 1000 ml. polyethylene bottles, care being taken to fill and cap the bottles so that as little air as possible was introduced with the sample. This precaution was to postpone any large scale atmospheric mixing until opened in the laboratory for filtration. Samples were filtered through a variety of Whatman filter papers. A preliminary filtration took place through Whatman If 5 filter paper, primarily to remove floating and coarse suspended matter. A second filtration through Whatman # 42 or # 44 filter paper was effective for all particles five microns in diameter or greater. Filters were mounted in glass funnels. Sediment samples taken from Burke Lake and Topsail Island were collected in 250 m1. wide mouthed, thick walled, polyethylene containers with special tight sealing lids. This material was returned to the laboratory, and the interstitial and associated water was 22 23 extracted by vacuum filtration through a Biickner funnel containing 2 - 5 micron teflon filters. The extraction process was aided by applying external pressure to the top of the sediment during filtration. The purpose of filtration was to remove all suspended matter, especially clays, since certain clays undergo cation exchange reactions with most of the dominant cations. To insure that the filtration was effective in removal of clays, an aliquot of Red Cedar River water was evaporated to dryness, and a portion of the residue was subjected to an infrared absorption analysis, by Mr. A. Bailey. The resulting spectra (Figure 4) indicates that no clay minerals were present; but it does indicate a strong concentration of carbonate and the presence of a variety of organic compounds. Following filtration, each sample was split into several ali- quots. A 250 ml. aliquot was acidified with concentrated HCl to a pH between 2 and 3, and then analyzed by A.A. for calcium. The acidification insured that the calcium would remain in solution and it aided in the maintenance of proper flow through the burner, insuring a more accurate and precise analysis according to the procedures of Bentley, and Lee, (1967). A Perkin Elmer Model #303 Atomic Absorption Spectropho- tometer, with a Perkin Elmer Ca - Mg Lamp was used in the analysis. The insturmental techniques described in the Perkin Elmer manual were followed for machine operation. Standard solutions were prepared in accordance with the above instrumental techniques from analytical 24 UOISSNUSUEIL %> 00» oNDO - sac N sow. go o sow a son: . 00¢ ODUHmom Amt/Hm HNOOU pom .Ho 002 P d H 80 .mnuoonwm OOHMHHCH “Hy mudenH oomH b . oomH . d 25 grade CaCO3, dilute HCl, and deionized water. Working curves were drawn plotting absorption versus calcium concentration over the ranges from 0 - 1 ppm calcium, in accordance with techniques described in the manual. The acidified samples were diluted 10 to 20 times with deionized water using pipets and volumetric flasks. Water used in dilution of samples and preparation of standard solutions was first distilled and then passed through a Barnstead Bantum Dernineralizer utilizing a mixed bed ion exchange column. Prior to sample analysis, at least five standard 10 ppm and 0 ppm calcium solutions were run to insure the instrument was prop- erly "warmed up", and giving stable repeatable absorption values on these solutions. Samples were then analyzed in groups of three, with 10 ppm standard solutions run for reference after every third sample analysis. Absorption values were converted to ppm calcium by refer- ence to the working curve, and multiplied by the proper dilution factor for the total calcium concentration. This concentration has been termed total, because the absorption process evaluates all the calcium present as previously explained. At least three non-acidified aliquots were analyzed for both free calcium ion and pH. Determination of calcium ion activity was made with an Orion Calcium Activity Electrode, model #92-90, in conjunction with a Beckman calomel reference electrode and an Orion model #401 Specific Ion Meter. Basic instrumental procedures were 26 followed in accordance with those recommended in the Orion Instruc- tion Manuals. A l x 10'1 M CaClz standard solution was prepared from ana- lytical grade CaClz and deionized water. This was diluted as required to make 1 x 10"3 and 1 x 10'2 M CaClz standard solutions, which were then used to standardize the specific ion electrodes. Standardization was complicated by the fact that individual ion activity is seldom equal to the total concentration of that ion in solution due to the effect of the activity coefficient. The activity coefficient, which is a function of the ionic strength of the solution, affects ion activity according to the relationship: 0' Ca2+ : a[Ca2+ MCa2+ where a Ca2+ = calcium ion activity X (332+ MCa2+ = total calcium concentration, in Moles per liter. calcium ion activity coefficient From calculations of the ionic strengths of the sampling areas (Table l), the calcium activity coefficients for these waters were extrapolated from a graph formulated by Garrels and Christ, (1965, p. 63). The calcium activity coefficient is always less than unity in the sampled sites, which means that the calcium activity will always be less than the total concentration. Since the ionic strength of a particular sample environment is essentially constant, the activity coefficients for that environment can be treated as a constant. Thus, 27 in an inorganic solution, where no other interactions with calcium ion are possible, simple subtraction of the effect of the activity coefficient should make the calcium ion activity readings equal to the total concen— tration, all other factors being equal. This subtraction was incorpo- rated into the standardization of the calcium activity electrode, so that its readings of calcium ion activity should be equal to the calcium con- centrations given by A. A. for an inorganic standard solution. Awareness of this inorganic activity coefficient was most important, because reactions of dissolved organics which would com- plex calcium ions would also result in calcium activity reductions. By removing the reduction caused by the inorganic activity coefficient, any calcium activity variations caused by interactions between organic complexes and this ion became readily apparent. Analysis of the l x 10'3 M CaClZ standard solution by A. A. gave a total calcium concentration of 36 ppm. To remove the effect of the activity coefficient, the calcium ion electrode was standardized so that the reading obtained for this standard solution was equal to an activity of 36 ppm. This activity corresponded to a reading of 0. 00 millivolts (mv. ) on the expanded mv. scale of the Orion Specific Ion Meter. Using the 0. 00 mv. setting and various strengths of known standard solutions, a working curve was drawn by plotting mv. versus calcium concentration. A reference chart which directly converted mv. t0 ppm calcium was calculated from this curve for intervals of 0. 25 mv. over the ranges expected for calcium concentrations in fresh 28 water environments (Appendix B). Samples were analyzed in either 100 ml. or 250 ml. beakers, depending on the size of the aliquot. Samples were stirred for at least fifteen seconds with a teflon stirring rod after insertion of the elec- trodes. A waiting period of at least two minutes was required for equilibrium to be achieved between the sample and the activity electrode. The operating characteristics of the calcium electrode were extremely troublesome. The presence of small air bubbles on the electrode membrane caused erratic and unpredictable behavior. This was manifested as rapid fluctuations in the meter readings. These bubbles were difficult to dislodge and often the only recourse was to extract all moisture from the outer surface of the membrane with absorbent toweling. This, in turn, would require restandardization. Further problems yielding erratic readings were caused by the slow and unpredictable decay of the liquid ion exchanger within the electrode body. This decay always led to high values for the calcium activity. This problem was detected when the values given by the cal- cium electrode exceeded those given by A. A. by more than that pre- dicted by the analysis of variance. Those experimental values of cal- cium activity which exceeded the total concentration by more than 5 ppm were invalidated. Errors caused by this decay were usually between 10 and 30 percent before detection. Experience showed that this part of the calcium ion electrode operated properly for as long as four months, and as short as one month. 29 Other analytical problems were created by particle contamina- tion of the membrane. This would set up wild fluctuations in the meter readings, often sending the needle off scale. This problem was recti- fied by replacement of the membrane. Prior to sample analysis, the activity electrode was checked for proper standardization using the 1 x 10'3 M CaClZ and 1 x 10‘2 M CaClZ standard solutions. The electrode was also checked with the l x 10.3 M CaClZ solution after every three sample analyses. Between each analysis the bodies of the activity electrode and calomel electrode were wiped dry with absorbent toweling to prevent contamination of the following analysis. The Orion model #92-32 divalent cation activity electrode was standardized following the procedures described for the calcium activ- ity electrode. The divalent electrode differs in that it responds to the combined activities of all divalent cations in solution; however, this response is not uniform for all divalent cations. Since the activity coefficients for all of the divalent cations present could not be deter- mined, and the electrode response was variable, the divalent activity electrode was not used as a precise measure of absolute proportions of complexes divalent cations present. Instead, the electrode was used to monitor variations in activity, since the response variations and activ- ity coefficients could be treat-ed as a constant in a given sample. The same analytical problems which plagued the calcium ion electrode also affected this electrode. 30 Concomitant with determination of activity, the pH of the sam- ple was measured using a Photovolt Digicord pH Meter, and their pH and calomel electrodes in accordance with the Digicord instruction manual. The Digicord was standardized in Beckman pH 7. 00 buffer solution at room temperature prior to each period of sample analysis. Drift from the standardized pH over a time period of four hours was negligible, thus restandardization after every three sample analyses was not necessary. The pH values were read when the digital readout stabilized to variations of i 0. 02 pH units. Since the nature of organic - inorganic interactions has not been previously investigated, there was a lack Of information on sam- ple deterioration with respect to degradation of the dissolved organic matter by atmospheric oxidation or other destructive processes. An experiment was run to examine degradation processes as reflected by release of calcium ions from organic complexes with time from sampling. The results of this experiment (Table 3) indicated that little calcium was released in the first 24 hours, but after this time calcium is released in large quantities. The geochemical implications of this time degradation are discussed in the next chapter. Analytically it means that natural samples have to be analyzed within 24 hours of col- lection, or valuable information might be lost. The fact that calcium is released with time presented another means of sample analysis to Obtain prOportions of complexed ions to free ions. 31 Table 3 Time Decay Experiment Red Cedar River Sample Calcium Hours Concentration in ppm 0. 0 77. 5 0. 5 77. 5 1. 0 77. 5 1. 5 77. 5 2. 0 77. 5 4. 0 77. 5 6. O 77. 5 10. 0 79. 0 24. 0 81. 0 34. 0 85. 0 48. 0 89. 0 72. 0 97. 5 96. 0 104. 0 120. 0 106. 5 148. 0 104. 0 Total calcium concentration by A. A. = 106. 8 ppm. 32 Analysis of marine and brackish water environments presented some special problems. High ionic strengths cause large depressions in the calcium ion activity (Garrels and Christ, 1965, p. 63). In addi- tion, high magnesium ion concentrations (at least three times calcium concentrations) cause interferences with calcium activity readings with the calcium ion electrode. Furthermore, extensive ion pairs and pos- sibly other unknown inorganic complexes were certainly present to con- fuse the data (Garrels and Christ, 1965, pp. 106 - 108). Thus precise results analogous to fresh water data were impossible to obtain. Fortunately, the characteristics of the complexes observed in previous fresh water studies supplied a means of investigation. Com- plexes are unstable with respect to time, releasing their bound cal- cium during degradation. Thus, by monitoring the changes in calcium ion activity within a sample for at least 100 hours, a good estimate can be made of the amount of organic complexification. In this way prob- lems involved with magnesium interferences, inorganic complexes, ion pairs, and ionic strength could all be treated as a constant, because these features of non-ideal inorganic solutions are not time dependent. A standard solution* was prepared for the calcium ion activity electrode which approximated marine and brackish water ionic strengths. The total calcium concentration was determined by A. A. to be equal to *standard solution: prepared by dissolving: 7. 00 grams of NaCl, 0. 50 grams of MgSO4° 7 H20, and 0. 25 grams of KCl in 56 m1. of 1 x 10"2 M CaClZ standard solution and bringing the volume to 500 ml. with deionized water. 33 333 ppm. This corresponded to a reading of +24. 5 mv. on the specific ion meter; thus, for these experiments the calcium electrode was standardized at +24. 5 mv. being equal to 333 ppm calcium. Utilizing the described instrumental techniques, measurement of the important geochemical parameters and their interaction with calcium ion activity was carried out. By perturbing the samples through addition of acids, bases, oxidizing agents, and other cations, the effects of these disturbances on calcium ion activity were investi- gated. All sample analyses were made within thirty-six hours of col- lection to minimize loss of complexed calcium caused by time degradation of the dissolved organic matter. RESULTS Introduction Results indicate that, in general, the presence of organic com- plexes of calcium causes the calcium ion activity coefficients to behave as a complex function of pH, Eh, and ionic strength. These results show that organic complexes play a significant role in many aqueous geochemical systems, and that the precise nature of this role is not a priori predictable. The results in the following sections present a number of different, but related aspects and conclusions concerning the relative and absolute magnitude of organic complexification of calcium ion; and the nature of the interactions between dissolved organic complexes and the inorganic aqueous geochemical system. Relative a_ric_1 Absolute Magnitudes giOrganic Complexification Fresh Water Environments: Relative proportions and absolute amounts of complexed cal- cium were determined over a six month period for fresh water envi- ronments in central Michigan (Figure l), with pH readings concurrently taken. Results of this investigation (Table 4) indicate that amounts of 34 35 0N.w 00 .H. mod 00¢ 0w.> >0.w Hm.» mng 0mg. mvH .w Hm.w 0N.w 0H .N. wm.m brim mm.w wm.w wm.w 0m.w 0m.m In Sam M sec m fimN Sam N NEH chow NNH o\sNN $0H chum o\sNH cam H oxom H NNN SEN ass N cam N cam N exam m sac m NNN HOONOHAHEOO «GoononH om Hym om 0H HH mH 0H NH NH HVH 0N HuN 0N mN Nm 0m NM VN EanH GH ooQOHOHHHQ nymphs? Hmonh CH GOENOHHHNOHQEOO EsHonO .HO OOSHHQmmz Hm mm hm mm mm we m0 Nm Nm hm Ho mm ow mo Hi. mm mm Nw w» Hub H; H0 Ema flH sas>asunw GOH ESHOHMU 00H mm 00 mm mm Nm 2. >0 N0 00 m0H mo moH moo m0H HOH >0H >0H 0HH 0HH 00H mm 825 CH GOH EsHonO HSOH a. 2an oe\o~\~ as\o~\m os\oz\~ oo\wa\~ oo\m\~ os\z\~ os\©~\H os\s~\H ms\m~\a os\am\a mo\o~\aa ws\oz\aa wo\NH\HH wo\wa\oa ms\sa\oa ws\o\0a ws\w\oL ms\w\oL mo\~\oH ms\OM\o 33H; wo\v\o onnH Ho>H~H umHoOO pom 380 gm assesses. pot/Hm pmHoOO pom GOHHmOOH OHQEmm u»:- I 36 0H .0 m5.5 00.5 00.5 mH.w 00.5 50.5 00.5 50.5 00.5 00.5 HH .0 50.5 mH .0 00.5 vH .0 00.0 Hm.w 00.0 In smm 053 see ssm ssm sow $0 OSH H gmH $mH Q00 NEH fiNH fiNm §NN OoonnHEOO unoonom mHu Nm HHV 0m 00 Nwmm00sw>asunw GOH EdHonO Hum 55 5w Nw mm H0 m0 m0 m0 mm m0 0m 0H» 0m 5m mm mm 50 50 8mm CH GOH EdHonO Hfios Lenses 4 £an 00\w\N 00\w\N 00\w\N 00\w\N 00\w\N 00\w\N ms\e\mz ws\s\NH wo\e\~z w0\e\ma wo\e\ma we\s\ma w0\¢\NH ws\v\~H ws\s\~H w0\w\ma m0\mH\s os\s~\~ oo\m~\~ 3mm Hmmvmxo $52 934 335m 503 0305303 non—m3 OvawH Oxpsm wchGme OvawH .HOZMH anuOU pom GOHHmOOH OHnHEmm 37 complexed calcium varied widely, both between different sampling sites and between samples taken at the same sampling site on different dates. Proportions of complexed and total calcium in the Red Cedar River varied from 5% to 40%, and concomitantly the amount of com- plexed calcium varied from a minimum of 1 ppm to a maximum of 39 ppm. The low values represent samples collected in January during a winter thaw, when the river was at flood stage. High values occurred in September, 1968, and February, 1969, when the river was at a low stage. These variations were probably influenced pri- marily by the rate of discharge; that is, low values probably represent dilution of a relatively constant amount of organic complexes. Surface waters in Lake Lansing exhibited much lower levels Of complexed calcium (8% - 15%), compared to the Red Cedar River. These values reflect much lower total calcium concentrations, about 38 ppm, and a smaller quantity of organic matter, thus leading to lower amounts of complexed calcium, only 3 to 6 ppm. In comparison to the Red Cedar River and Lake Lansing, sur- face waters in Burke Lake exhibited virtually no complexed calcium. The water in Burke Lake is primarily derived from lake bottom springs which probably contain little organic matter. This factor plus significant rates of influx and outflow, and the relative youth of the lake (an old gravel pit) is reflected by the early stage of eutrophication that presently exists. This is probably why little 38 or no organic complexing was observed in these surface waters. Interstitial waters from dark colored organic rich sediments exhibited enhanced levels of calcium complexification. About 45% of the total calcium was complexed in pore waters extracted from the upper five centimeters of Burke Lake sediments. Total calcium ranged from 77 to 94 ppm, with from 30 to 48 ppm calcium existing in a complexed state. In comparison, the free surface waters in Burke Lake ranged from 63 to 83 ppm total calcium with only 0 to 8ppm bound in complexes. Thus even when organic complexification plays only a slight role in overlying waters, it is an extremely important fac- tor in underlying sediment interstices, a geologically important envi- ronment. Organic complexes in pore waters can act as agents for solution or precipitation. Cations may be sequestered by complexing agents in sediment pores, causing solution of available minerals, even when activity is at saturation. Reactions involving the complexing agent and calcium ion can cause mobilization or precipitation of calcium ion under circumstances regarded as inirnical by the inorganic model. Thus sediment pores can serve as a continuous source of cations as long as a supply of complexing agents is available. 39 Marine and Brackish Water Environments: The aforementioned results indicate the importance and wide- spread occurrence of dissolved calcium complexes in fresh water environments. To determine whether this was strictly a fresh water phenomenon or instead was a general effect, a reconnaissance study on free and sediment pore waters from marine and brackish water environments was undertaken near Topsail Island, North Carolina (Figure 2). The North Carolina results (Table 5) show significant increases in calcium ion activity between initial readings taken at the time of sampling and determinations after five days. The difference between these two readings, or the amount of calcium released, is here taken to represent complexed calcium. The proportion of complexed cal- cium is calculated by dividing the amount released by the total amount, which is assumed to be equivalent to the amount present after five days. Marsh water samples from Topsail Island had a normal marine ionic strength. Total calcium present was 420 ppm, of which 60 ppm initially were complexed, or about 14% of the total calcium. A surface water sample taken near Thomas Landing, was from a brackish lagoon, which had a distinctly yellow -brown color. Shapiro (1964) found that yellow - brown coloring matter in natural waters was a good complexing agent for iron and other metals. The results of this 40 «EN Hi» NNN NmH .8025 AN HCOEHHOOm Goowme mGHOGwH mmEOaHH. AH chNH HuN 00N NwH gene? Cw unoaHpom coomme M5934 mgoHHH H .52 mm SN mm: A825 .0 Eossemm soommq 3853 380.? o o0mH ow wON 00H HBO? ow HGOEHHOOm noommd UGmHmH HHmmmOH O exemH wN 00H m5H nous? GoowmIH maHHuQmiH mmEOHHH. m o\sHuH 00 0N0 00m pooh? Hmnmz HummHmH HHmmnHoH < Ema RH 8an m new?» CST/304% 0333800 0833800 35.33% HHOH Co.DHonU GOHHmOOwH 395mm HavonOnH GOH SSHOHMO aoH EDHOHCO HmHHHnH mugs? smHvHOmum Hone main: GH GOHHmOHHHNOHnHEOO EDHOHmO .HO @0503me m OHQmH 41 study tend to confirm his findings because 28 of the 198 ppm calcium were found to be complexed in these colored waters. As in fresh waters, the largest amount of complexed calcium was found in pore waters which, in this instance, were extracted from black, organic rich, sediment from the vicinity of Thomas Landing. This sample contained 238 ppm total calcium, of which 44 ppm or 24% was initially complexed. Generally, these results indicate that marine and brackish pore waters associated with dark - colored, organic rich sediments were enriched two or three times in complexed calcium with respect to the overlying free waters. The magnitude of organic complexes of cal- cium, precisely determined in fresh water, is shown to be present in the same relative proportion in similar marine environments. We must conclude that although the increased ionic strengths of marine and brackish waters makes them chemically different in some respects to fresh water, this is not a major factor in controlling the presence of organic complexes. The presence of excess amounts of cations other than calcium in marine waters probably is responsible for the slightly lower proportions of observed complexed calcium. If this is so, then according to results described in Cation Exchange, total amounts of organically complexed cations might be equal to or greater than in fresh waters. 42 Time Degradation Having determined absolute and relative magnitudes of calcium complexification in a range of natural environments, the nature of the complexing agent was deduced by observations and interpretations of changes in the calcium activity and pH in controlled experiments. Since it was shown (Table 3) that calcium activity in samples was variable with time, a more thorough examination of time - organic matter interactions served as a starting point. Figure 5 shows the change in calcium ion activity with time expressed as change in the proportion of initially bound calcium released versus time in hours. In the first ten hours after collection no breakdown of calcium complexes was' observed. This could repre- sent a period in which less stable, non-complexing dissolved organic matter is being broken down; or it could represent a time period in which the sample does not degrade. Between ten and one hundred hours, the degradation process proceeds in a linear manner. Degradation probably results from actions of two factors, bacterial action or atmospheric oxidation, either singly or in combi- nation. Although the present investigation cannot clearly determine the degradation mechanism, the linear nature of the curve in Figure 5 does indicate that the rate of degradation is more or less constant. The last 5% to 10% of the bound calcium was released more slowly than the previous portion as evidenced by a change in slope. 43 100% T . O 80% "F O "o a) g 60% "' ,2 o u 52 .9. E .2 '3 40% u o m o 'o o X ,S’. o. E 0 8 “a 20% A o no 53 . £1 a) B O o D. h- : : 4 ‘r s 24 48 72 96 120 HOURS Figure 5: Time Degradation Effects on Complexed Calcium Ions 44 A reason for this might be the presence of a group of complexing agents which are distinctly more difficult to degrade; or that the general degradation process floods the rnicro-environment with products that retard further degradation. The results of this experiment have important geologic impli- cations. No currently known aqueous geochemical model can explain the time degradation trend. Explanation of these results must lie in considering the active presence of dissolved organic complexing. material. In addition the calcium ion activity increase in time indi- cates that the calcium is bound to an organic host that is relatively easily degradable. Since these organic complexes are unstable, they must represent a significantly different type of chemical species than the well-known, highly stable, porphyrin complexes such as chlorophyll. Another time degradation experiment was performed in which pH was monitored concomitantly with calcium ion activity (Table 6). These data indicated that the pH increased with time. Figure 6 was plotted to examine the nature of the relationship between calcium re- lease and changes in pH. This plot indicates that the increase in pH occurs very quickly compared to calcium release, and is essentially completed before the majority of the calcium complexes break down. These pH changes are apparently due to a set of reactions independent of organic complex degradation. It is interesting to note that the major increase in pH occurs in the early portion of the calc1um - orgamc 45 - Table 6 Time Degradation of a Red Cedar River Sample Time Calcium Total Complexed in Activity Calcium Calcium pH Hours in ppm by A. A. O 65 97 32 8. 32 2 68 97 29 8. 39 6 68 97 29 8. 52 24 75 97 22 8. 64 48 89 97 8 8. 55 complex degradation curve. This means that these pH trends might be related to breakdown of other less stable, non-complexing types of dissolved organic matter. The unstable organic matter may be high molecular weight molecules which are easily broken down by a variety of processes, and it is not uncommon that this decay removes hydro- nium ions from solution by hydrolysis (Degens, 1965, pp. 207, 209, 226). This degradation might in fact provide a reservoir of new com- plexing agents which would retard calcium ion activity increases. Pe rtu rbation Expe riments Introduction: Because ionic strength, pH, and Eh are important parameters for all geochemical models, natural samples were perturbed in a con- trolled manner to evaluate the response of organically complexed cal- cium to changes in concentrations of other cations, pH, and Eh. 46 24 hours 8. 6O "' 48 hours 8. 50 ‘* pH 8. 40 It i '2 hours 4, : 1L % _: 20% 40% 60% 80% 100% % of Bound Calcium Released Figure 6: Release of Complexed Calcium vs. pH with Time 47 Cation Exchange: Detailed studies of synthetic organic complexing agents has shown that such molecules have definite orders of preference for types of cations with which to complex (Martell and Calvin, 1952). This preference is a function of the size, charge, and shape of the candidate cation. If cations are introduced into a system containing a complexing agent which has previously bonded to a less desirable cation, an ion exchange reaction will take place in which the more preferred cation will be sequestered and the previously bonded cation released. The kinetics of the ion exchange are increased if the new cation is intro- duced in excess. To test whether or not calcium was the most stably complexed cation, excess amounts of Mg2+, Fe2+, Fe3+, Ni2+, Cu2+, and Ba2+ were added to separate aliquots of Red Cedar River samples. With the exception of Ba2+, all of these cations caused release of bound calcium ion. The results of these exchange experiments are shown in Figure 7 (data in Appendix C). The greater relative size of the barium ion is probably the reason it did not effect calcium release. The Fez+ effected more calcium release than Fe3+*, which is in accordance with ligand field theory predictions. >kOrganic complexes of iron are especially important because this ele- ment is geologically common and the geologic record attests to its mobility in the past; yet it is almost insoluble inorganically in most common natural water. Thus, organic complexes sup- ply a possible mechanism for the transport of iron in neutral to alkaline solutions. CCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCCC IIIII RRRRRR F83... "192+ O \\\\B 49 Addition of copper caused the greatest calcium release. This high release is apparently not predicted by considerations of size and charge, because Fe3+ should be more acceptable due to its smaller size and greater positive charge. Thus, the greater acceptability of copper must be due to highly specific steric preferences, character- istic of these complexing agents, which is predicted by considerations based on ligand field theory. The low values of exchange for Ni2+ are not in accordance with predictions based on this theory. Reasons for this lack of exchange merit future investigation. Perhaps most important was the result showing that excess magnesium ion is preferentially complexed over calcium. In marine environments where magnesium ions have a 3 to 1 ratio over calcium ions, the amount of complexed magnesium should exceed that of com- plexed calcium. Results of the North Carolina study indicated that about 15% of the magnesium ions may be complexed in those environ- ments. This could have important ramifications on carbonate equi- libria in marine and brackish water environments where large amounts of magnesium ions may exist in organic complexes, especially in sedi- ment pore waters. Release of this bound magnesium, with lesser amounts of calcium when the complex degraded could cause precipita- tion of carbonate cement, and possibly act as another mechanism for dolomitization. Summarizing these results, it appears that ion complexification is a function of the availability of an individual cation at the time of 50 complex formation, and the concentration of that cation with respect to others in solution. The fact that exchange took place shows that those divalent cations smaller than calcium are transported in a complexed state. The role of organic complexes in weathering takes on new impli- cations, because slightly soluble cations such as iron and copper can now be effectively solubilized and transported as organic complexes. 1H Pe rtu rbations: The previous experiments have shown that the amount of ion exchange increased as calcium complexes were presented with pro— gressively smaller cations. Carrying this experimental approach one step further, various concentrations of hydronium ions (H+) were added to various samples, and the reactions monitored for changes in the calcium ion activity, divalent ion activity, and pH. As a preliminary experiment, samples taken from Rose Lake and Lake Lansing were titrated with 2 x 10'2 M HCl and monitored for resulting pH. A region of extensive buffering was observed between pH 7. 6 and 6. 0 (Figure 8). Because buffering in the inorganic model is almost entirely attributed to bicarbonate ion, a synthetic inorganic Rose Lake solution* was prepared and similarly perturbed to evaluate magnitudes of purely inorganic buffering. Samples of this synthetic *Synthetic Rose Lake solution contained sodium bicarbonate, calcium chloride, and magnesium sulfate in concentrations closely approximating those observed in natural samples (Table l). 51 .01 ENod yo .2: n m _ . . q . 9m .. oé fixammy awkfimmom 2352...... -on 1 .33 west K 83.023 .33.!“ . o.o .. o.» / .. o.m mmx' pH '-. INITIAL SAMPLE pH-s. 8.36 0.00 7.00 - \ SYNTHETIC RED CEDAR RIVER SAMPLE 6.00 '- 54 exchange reaction between hydronium ions and organic complexes similar to that commonly observed with artificial complexing agents such as EDTA, and oxalic acid. A synthetic Red Cedar River solution* was prepared to test if the calcium ion activity increase upon addition of hydronium ion was due to the presence of inorganic complexes. This solution was then titrated with 2 x 10“2 M HCl following procedures identical to those used with natural samples. The acid titrations of the synthetic Red Cedar River sample caused no changes in the calcium ion activity (ppm); and as in preliminary experiment, no buffering region was observed (Figure 9). Thus, increases in calcium activity in the natural sample must be due to interactions between added hydronium ions and organic complexes of calcium ions. The additions of base (2 x 10'2 M KOH) caused a slight lowering of calcium ion activity due to dilution until the pH increased to the point where carbonate precipi- tation began, which resulted in more substantial calcium activity decreases. Part of the reason for calcium release upon addition of H+ is probably due to ion exchange between the very small H+ ions and pre- 2+ viously bound Ca ions. Because the exchange process is based on charge balance theory, two H+ ions would have to be complexed for *Synthetic Red Cedar River solution was prepared by dissolving 0. 37 grams of MgSO4' 7HZO; O. 25 grams CaCO3; 0. 083 grams of NaHCO3 in a liter of deionized water. 55 every Ca2+ ion released. Calculations based on data in Appendix D indicate that approximately 3. 5 to 5. 0 H+ ions were added for each Ca2+ ion released to pH 6. 7, because further additions of hydronium ions caused little or no calcium release below this pH. Because more than enough H+ ions were added to account for the observed amount of calcium released and pH changes, it must be concluded that many of the added H+ ions are playing another role. This fact was implied in discussions of the time effect on pH and cal- cium ion activity in a previous section. The majority of the added H+ ions were probably involved in degradation processes (hydrolysis) with the dissolved organic matter. The need for, and role of, H+ ions in hydrolysis of organic matter is well-known (Degens, 1965, pp. 209, 226, 254). An important aspect of the cation exchange processes is reversibility. To test for reversibility a number of experiments were conducted in which a Red Cedar River sample was first subjected to an acid titration, then a base titration was performed, then re-acidifi- cation. The results of this are tabulated in Appendix D and summarized by Figure 10. The results indicate that the interaction between hydro- nium ions and the organic complexes is not a simple exchange process, for no reversibility was achieved. Additions of acid brought increases in the calcium ion activity until all complexed calcium was released. Intermediate addition of base caused no change in the calcium ion activity. pH on" ACTIVITY A$ppm 8.5 8.3 i- 8.2 8.! 8.0 1_1 7.9- 7.8k- 7.7 - 7.6.- 7.5L- 7.4b 96- 94L- 92- 90- aer- 86L 84b 02- 80 56 Figure /0 CALCIUM ACTIVITY-pH RELATIONSHIPS ACID ml. of 2xIO'2 M HCI or KOH 57 Generally, these results indicate that organic complexes are quite sensitive to additions of hydronium ions. Acidification of an environment will be buffered by the presence of organic matter. Part of the buffering effect is probably due to preferential release of cal- cium ions for hydronium ions in the complexes, and the other part is probably due to hydrolysis of high molecular weight organic matter. These results indicate that a large portion of the complexes are involved in degradation processes as well as cation exchanges, because acidification irreversibly alters part of the dissolved organic spectrum. Oxidation Perturbation: Previous experiments investigated interactions between organic complexes of calcium ions and time, variations in concentrations of other cations, and variations in hydronium ion content. The following experiments use similar techniques to investigate the effect on organic complexes of oxidation potential changes. The previously cited time decay study indicated that there is probably an oxidation effect on the organic complexes. To test the effect directly, six aliquots of Red Cedar River water were divided into two groups of three aliquots each. To one group enough 30% hydrogen peroxide (H202) was added to result in a 2% concentration of peroxide in the aliquots. No peroxide was added to the other group. Calcium ion activity and pH was monitored in these groups for two days. 58 The results (Appendix E) of this experiment show an important difference between the behavior of the treated and untreated sample with respect to calcium ion activity (Figure 11). Results indicate that this slightly enhanced oxidation potential has actually created new com- plexes which lowered the calcium ion activity. This oxidation may be breaking down high molecular weight organic material and creating some relatively simple, dicarboxylic acid - like molecules which are doing the complexing. Conditions causing general complex degradation took hold in both treated and untreated groups within ten hours, and the previously observed linear degradation trends were produced. However, those samples treated with peroxide always contained relatively more complexed calcium at a given time. It appears from these results that weak oxidation processes, such as atmospheric oxidation, actually create complexing species from available organic material. This may mean that when environ- ments change from reducing to oxidizing, additional organic complexes are created which can act on the available cations. The effects of weak oxidation are constructive, whereas the effects of strong oxida- tion (Chapter 2) were shown to be destructive with respect to cation complexes. Since addition of peroxide, a strong oxidizing agent, caused decreases rather than increases in calcium release, one must conclude that the controlling degradation factor is something other than simple oxidation. The observed pH variations in this experiment, involving rapid 59 100 1' bn— ._ o —omn—0_A.A. 00 \O O O l I Calcium Ion Activity in ppm \l o L 0 1 1 1 I j r 10 20 30 40 50 HOURS Figure ll: Calcium Activity - Eh Relationships 60 increases in pH with time, were similar to those previously reported, both in peroxide treated and untreated samples. Slight depression of pH in the peroxide treated sample was due to the fact that Eh is not completely independent of pH. Increase in Eh will cause decreases in the pH, and vice versa (Garrels and Christ, 1965, p. 137). The previously discussed experiment investigated the relation- ship between increased oxidation potentials and the resulting changes in calcium ion activities and pH. An experiment was especially designed to test for Eh - pH interactions; that is, if induced changes in these two parameters resulted in calcium ion activity levels that were not simply additive. The following table summarized this design. Samples in Table 7 were analyzed over a two day period, and results are summarized in Appendix E. Table 7 Eh - pH Interaction De sign No HCl* 2% Hc1* 5% 140* 10% Hc1* added by volume by volume by volume No H202” 3 3 3 3 added replicates replicates replicates replicates 2% H202M 3 3 3 3 by volume replicates replicates replicates replicates * 2 x10'2 M HCl 3 0% H2 02 HCl **H202 II 61 The results of this experiment indicate that no interaction was present between additions of acid and peroxide. This lack of interaction is shown in Figure 12 which plots calcium activity against time. This graph shows that additions of acid did not alter the basic shape of the peroxide degradation curve with respect to the non - peroxide treated curves. The acid additions simply acted in an additive fashion increasing calcium ion activity at any given time. In all samples treated with peroxide, the calcium activity was lowered with respect to untreated samples. Another clear result is the observed presence of a pH limit between 8. 5 and 8. 7, over time in the Red Cedar River (Figure 13). These observed pH limits are probably related to the quantity of hydronium ions needed for hydrolysis reactions with available organic matter, coupled with any use of hydronium ions in other types of degradation processes such as bacterial metabolic reactions. Since additions of peroxide did not affect observed pH, it appears that those reactions causing changes in observed pH are not the result of simple oxidation. In summary, this experiment showed that no interaction existed between concomitant additions of peroxide and acid in natural samples. That is, that calcium ion activity in samples treated with peroxide displayed lower levels than untreated ones over the entire recorded Span. Additions of acid uniformly raised the curve. In no sense was a relationship observed where combinations of peroxide and acid levels 62 Figure [2 Eh-pH EFFECTS ON CALCIUM ACTIVITY l00 " 00‘ CALCIUM ACTIVITY IN ppm 63 Figure I3 Eh-ACID EFFECTS ON OBSERVED pH -— No H2 02 added -- -2 ”/9 H2 02 by volumn 5 0/0 HC/ 7.0?- » I0 % Hc/ 6.5L L 4 1 1 I I0 20 30 4O 50 HOURS 64 produced abnormally low or high calcium activity results indicative of interaction. The previously reported lowering of calcium activities by increased oxidation potentials, and increases in calcium activity upon addition of acid were both evident in this experiment. Degradation of dissolved organic matter causes increases in observed pH. The degradation process probably consists largely of hydrolysis reactions, the limits of which may be controlled by the micro-biological ecology of the environments. Limiting pH values are probably a function of the total quantity and molecular types of available dissolved organic matter. Role of Organic Comflexes in Calcite Equilibria in Burke Lake Previous discussions have demonstrated the effectiveness of organic complexes as solubilizing agents, but they have not considered the role of complexes in precipitation reactions. To be an important factor in mineral precipitates, the amount of cation released upon complex degradation must cause supersaturation of that ion to initiate its precipitation. Many aqueous environments are near saturation, or indeed are saturated, as Burke Lake waters are with respect to calcium, and further cation additions should cause precipitation. Since a large portion of the total calcium in Burke Lake sediment pore waters is bound in complexes, this environment provided an ideal situation for a theoretical examination of the role of organic complexes in calcite equilibria. 65 Relevant equations defining the inorganic relationships are as follows (Garrels and Christ, 1965, p. 76): (Ca2+) (co32-) _ —8.3 Kcalcite ” I 10 CaCO3 calcite (11+) (H603') K : : 10-6°4 H2CO3 01+) (co3z‘) K _ = 10.103 HCO3 .2 HCO3' + _ (H ) (0H ) 44 0 KH : : 10 . 20 H20 H CO K ( 2 3) _ 104.47 C02 P _ C02 gas An equilibrium line expressing the solution or deposition of calcite can be drawn by solving the following equations and plotting the results as log Ca2+ versus pH. First a theoretical MCO32- can be obtained from (10'10' 3) (0 HCO3‘) (MHco3') (Q H+) (Y €032 -) — _. MCO32' and when calculated for the pH range 7. 5 to 8. 5, and using the experi- mentally determined HCO3- (Table 8), for pH = 7. 5, H+ : 10'7° 5 M - 10-5.09 (3032' _ —7 5 0 19 ' (10 ‘ ) (10' ‘ ) 66 Table 8 Calcite Equilibria Data for Burke Lake Total Calcium pH Calcium Activity Water Extracted from 6 Burke Lake Sediments #1 7. 64 10'2- 63 10'2- 8 #2 8.13 10-2-68 10-3~03 #3 7.46 10-2 70 10-2 84 #4 7.68 10-2-66 10-2-94 #5 7.73 10-2-72 10-2.95 #6 8.10 10~2-68 10-3-0O Burke Lake Surface \Vaters A. 7.66 10-Zo80 10-2-8O B 7.69 10-2-70 10-2 73 c 7.67 10-2-79 10-2 80 D 7.60 10-2-79 10-2-80 .E 7.67 10-2.79 10‘2-80 Ionic Strength 2 l. 2 x 10'2 M _ _ -2.43 MHCO3‘ — 225 ppm — lO _ -0.05 YHCO3"10 _ = 10.019 X6032 J’ca2+ = 10.019 I, 2 1035 C02 67 Then for H+ : 10-8' 5, : 10-4. 09. c032- Calculation of theoretical MCa2+ at equilibrium with a given pH involves: 10-8. 3 Mca2+ (Mco32-) (Xco3z-) (may) at pH = 7. 5 this expression yields: 10'8. 3 -2 83 M = = 10 ° C 2+ a (lo-5.09) (10'0.19) (lo-0.19) 10-3.33. and at pH = 8. 0 this yields MCaZ+ = This result is plotted in Figure 14 as calcite equilibrium line A. This equilibrium line can also be calculated by assuming that atmospheric PCOZ is fixed. Garrels and Christ (1965, p. 53) find that 10'3’ 4 moles of Ca2+ are in equilibrium with calcite at pH 8.4, assuming the calcium activity coefficient equal to one. Since the calcium activity coefficient in Burke Lake is equal to 10'0‘ 19, this -3. 5 . . equilibrium concentration becomes 10 9, at pH 8. 4 and Line B 1S its expression in Figure 14. These two methods give theoretical limits as to the position of the calcium ion - calcite equilibrium line. Table 8 lists the relevant experimental data for Burke Lake, and the data are plotted in Figure 14. The open circles are plots of Burke Lake surface waters, which indicate that the lake water is satu- rated with respect to calcite. The lower points of the vertical lines are the calcium ion activities taken from the specific ion electrodes for the pore waters. The tops represent the total calcium concentration 68 Figure I4 CALCITE EQUILIBRIA IN BURKE LAKE A 8 ° ’ Aca .. 0 one“ — --Amounf of Ca"L ion '0 - 2- 6 - bound in organic '1" ' complexes uncertainty 7 #4 3/0-3.eo I #6 #2 I I 7 7 I r5 l I I I l I I I0.2.8 _ | I l I l I l | | mCa“ I I I l I I I I l0.3.0 _ . I O l0.3.2 __ SOLUTION PRECIPITATION 1 1 l L 7.0 7.2 7.4 7.6 7.8 8.0 8.2 pH 69 measured by A. A. The length of the line indicates the amount of cal- cium contained in organic complexes. Organic complexes are playing an important role in this environment by acting as reservoirs for large amounts of calcium ions. The complexes can play two important parts. One, a sudden degradation of these complexes would flood these waters with great excesses of calcium ions, which should cause immediate precipitation of calcite, probably as a cement. The other role these complexes may play is as an intermediate stage in degradation of organic matter where they represent new complexes being created by oxidation processes. These new complexes may dissolve calcium from the sediment and act as a transport mechanism to the surface. The potential solution of calcium could actually create porosity within these recent sediments. These considerations leave no doubt that organic complexes play an important role in the carbonate equilibria in this environment. An important role for dissolved organic matter must be acknowledged if aqueous geochemical models are to be mean- ingful dynamic representatives of the natural system. NEW AQUEOUS GEOCHEMICAL MODEL Introduction A geochemical model is a set of factors that when integrated simulate the geochemical environment. This model can be used to predict geochemical behavior or infer genetic processes from geo- chemical products. Although any sort of listing of properties constitutes a model, the essence of a model is best expressed in mathematical terms. Results of this research identify a number of new factors that must be included in aqueous geochemical models. The presently accepted inorganic model expresses characteristics of the geochemical system as succinct ratios or equations, and therefore additional factors should be expressed in a similar manner, if only for comparative purposes. All aqueous geochemical models are centered around the con- cept of ion activity. In currently accepted models, ion activities are almost entirely independent of the other geochemical parameters, except at equilibrium. Relationships between ion activities and these 0 o o 0 a * . . . parameters are summarized in equihbrium constants . Equllibrium >I‘Equilibrium constant (K), for a reaction aA + bB 2 CC + dD, is CC + Dd : Aa + Bb’ 70 71 constants are simple ratios describing ion activities at the point of equilibrium between solution or precipitation of the solid phase. Variations in the important inorganic parameters, pH or Eh, will alter equilibrium relationships causing either solution or precipitation. When similar changes occur at other than the equilibrium point, they do not cause alterations in ion activity. Assuming that inorganic effects are paramount, the equilibrium constant is sufficient to adequately describe aqueous geochemical systems. Use of the present inorganic model results in significant overestimates of some cation activities, and grossly underrates the capability of natural waters to dissolve and transport ionic species. These insufficiencies are due to the previous assumptions that the ionic strength is the only control on ion activity except at equilibrium where it is controlled by solution or precipitation of its solid phase. In addition, equilibrium constants yield information on reaction direc- tion, but give little information of rates or ion activities when condi- tions are other than at equilibrium. It is probably due to this factor that mass balance estimates have so often yielded values bearing little resemblance to those predicted by observed products. Results of this investigation have shown that inorganic models are not correct. In addition to ion competitions existing at equilibrium between cation and solid, there is competition with other species in solution. The effect of organic complexing agents is very important in fresh waters, and compares well with effects of inorganic 72 competitions, such as ion pairs, in marine waters. In fact, it can be demonstrated that under conditions of saturation or supersaturation (predicted by the inorganic model), solution of the solid phase could occur. Results given in earlier sections indicated that calcium ion activity varies with increases in Eh and decreases in pH. Indeed, without perturbation, activity increased with time within a single sample. These variations were ascribed to interactions between calcium ions and degradable organic complexing agents. Example of the New Model Limitations of the data available herein make it impossible to formulate a general universally sufficient empirical model, but one can examine the nature of interrelationships of data from this study in terms of relative magnitudes, rates, and functional types. As an example of the characteristics of such a model, general functional relationships are presented for the Red Cedar River, utilizing the data in Appendix E. These expressions predict calcium activities as a function of acid additions and time from sampling. Although not explicitly reflected, it must be kept in mind that the presence of organic complexing agents is implicitly necessary for validity of the following relationships: Equations (1) to (3) express calcium ion activities as a function of time (T) with various amounts of acid added. 73 (1) a Ca2+ 2" 65 + 0.42 T -- no acid added (2)51 Ca2+ = 77 + 0. 36 T -- 5% by volume 2 x 10'2 M HCl (3) a (332+ = 83 + 0. 36 T -- 10% by volume 2 x10-2 M HCl where ¢Ca2+ represents calcium ion activity in moles per liter, and T represents time in hours after sample collection. Equation (4) expresses changes in calcium activity as a function of acid additions where ”x" is acid in percent by volume. (4) aCa2+ = 65+l.9x Combining these four equations (assuming slopes in (l), (2), and (3) are approximately identical) into a single expression gives (5)61 Ca2+ : 65+1.9x+0.38 T. Equation (5) represents a first attempt at a new model which takes into account variations in activity due to effects of organic complexing agents. Figure 15 illustrates the effects of equations (1), (2), and (3), as functions of calcium ion activity versus time. Solid dots are observed values. Open circles represent theoretical values calculated from equation (5). The linear nature of calcium ion activity increases can be clearly seen in this graph. The effects of acid addition do not appreciably alter either the linear nature or slopes of the calcium - organic complex degradation curves. The only effect of acid addition is one of increasing the amount of calcium released at a given time by a constant; that is, the amount of complexed calcium remaining is CALCIUM ACTIVITY IN ppm 74 Figure I5 OBSERVED ct:2+ ACTIVITIES vs. NEW MODEL |00 90 { =95 96 Confidence limites about each data point. 0 OBSERVED CALCIUM VALUES 0 CALCULATED CALCIUM VALUES 0 FROM NEW MODEL 75 lowered by a constant directly proportional to the volume of acid added. Another example of the type of supplement needed in the current model is expressed by the following set of equations derived from values taken from Figure 13, relating the effects of time on pH: (6) pH = (0.120 - 0. 0001328 T).1 for untreated sample (7) pH = (0.013 - 0.0000461 '1‘)"1 for 5% by volume 2 x10'2 M HQ (8) pH = (0.0122 - 0.000063 T)'1 for 10% by volume 2 x 10-2 M HCl Equation (9) expresses changes in pH as a function of acid additions where "x" is acid in percent by volume. (9) pH = 8.33 - 0.209 x The reciprocal linear equations (6), (7), and (8), in combination with equation (9) represent another example of the type of functional relationships needed to augment the present inorganic model. The geochemical meaning of the pH functions was previously discussed. In conclusion, the effects of organic complexing agents on cation activity in synthetic solutions is not a new concept (Martell and Calvin, 1952); but prior to now determination of the relative impor- tance of organic complexes on mineral equilibria was non-existent. The empirically derived models presented herein are only a first step toward generation of a much larger, general model for aqueous environments. It is encouraging that supplements to the inorganic model will 76 be of a simple mathematical form with relationships being quasi - linear or reciprocally linear. It is also obvious, that opposed to the functional simplicity of these effects, universal, generally acceptable models will have to be empirically determined from natural samples. This is in contrast to the inorganic model which was derived from and is universally correct for synthetic systems. Collection of these empirical data must involve determination of rates of influx and concentrations of mineral acids and bases, organic complexing agents, dissolved organic matter, and oxidizing and reducing agents. Added to this will be characterizations of the geochemical role of the micro-biologic community. This formidable task will require interdisciplinary unions of thought involving geolo- gists, ecologists, micro-biologists, and men of other scientific fields, to fully structure a relevant dynamic model of aqueous environments. SUMMARY AND CONCLUSIONS Organic complexes occur in significant amounts in many aqueous environments. These complexes often play an important role in mineral equilibria, especially in sediment pore waters. Rates of influx and degradation of dissolved organic matter certainly influ- ence and probably control solution and precipitation of minerals in many aqueous systems. Organic complexes cause cation activities to behave as complicated functions of Eh, pH, and the mass budget of dissolved organic matter. Interactions between dissolved organic matter, and degradation processes within samples caused increases in pH and calcium ion activities with time. The pH increases occurred quickly in the first few hours after sampling, whereas calcium ion activity increases occurred at a much slower, linear, rate over several days. Hydrolysis reactions affecting sample pH may have created products which tended to retard degradation of calcium- organic complexes. Artificial perturbations of natural complexes revealed aspects of the nature of the organic complexing agents. Ion exchange reactions were observed when smaller divalent and trivalent cations were added in excess. Preferences were in accordance with predictions by 77 78 current ligand field theory except for the anomalously low exchange with nickel ion. The observed exchanges were probably controlled by steric factors of the complexing agents. Artificial additions of acid caused release of complexed cal- cium, but additions of base caused no change in calcium ion activities. Interactions between added hydronium ions and calcium- organic com- plexes is irreversible. Additions of acid caused hydrolysis reactions which produced a strong pH buffering between pH 6. 6 and 7. 6. Slight increases in oxidation potential probably created new complexes which lowered calcium ion activities; but strong oxidation tended to destroy organic matter, releasing complexed cations. Observed oxidation interactions with organic complexes were different than observed time degradation reactions, indicating that degradation processes were controlled by other factors. Concomitant increases in Eh and decreases in pH created no apparent interactions. All observations indicated that decreases in pH caused increases in calcium ion activity, whereas the increases in Eh caused decreases in calcium activity. There were no anomalous differences observed in the variously treated samples. The significance of interactions between dissolved organic matter and the inorganic system suggests a modification of the present aqueous geochemical model. To demonstrate the nature of the inorganic - organic interac- tions, an empirical model was developed for the Red Cedar River from 79 observed data. This new model characterized calcium ion activities as a function of the magnitude of the complexed portion, changes in pH, and organic complex degradation rates with time from sampling. It is encouraging that this supplement to inorganic models will be of a mathematically simple form with relationships being quasi - linear or reciprocally linear. It is also obvious that opposed to the functional simplicity of these effects, universal generally acceptable models will have to be empirically determined from natural samples. This concept is in opposition to the inorganic model which was derived from and is universally correct for synthetic systems. Gathering empirical data must involve determination of rates of influx and concentrations of mineral acids and bases, organic acids, dissolved organic matter, oxidizing and reducing agents, and the roles of geochemically active micro-organisms. This formidable task should be the central effort of low temperature aqueous geochemistry. Interdisciplinary efforts will be required to precisely develop this new model. LIST OF REFERENCES LIST OF REFERENCES References Cited Baas Becking, L. G. M. , Kaplin, I. 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E. , 1965, "Studies on dissolved carbohydrate in Cape Cod Waters, 111 Seasonal variation in Oyster Pond and Wequaquet Lake, Mass.", Lim. 8: Ocean. , V-lO, #4, pp. 249-256. Wangersky, P. J. , 8: Gordon, D. C. Jr. , 1965, "Particulate carbonate, organic carbon, and Mn in the open ocean", Lim. 8: Ocean. , v-10, #4, pp. 544-550. Williams, P. M., 8: Zirino, A., 1964, "Scavenging of Dissolved Organic Matter from Sea Water with Hydrated Metal Ox1des", Nature, V-204, pp. 462-465. Wilson, R. F. , 1961, "Measurement of Organic Carbon in Sea Water", Limnol. 8: Ocean. , V-6, pp. 259-261. APPENDICES 87 APPENDIX A Preliminary Experiment Sample taken from Red Cedar River, split into six aliquots Total Calcium + Magnesium concentration measured by EDTA - Eriochrome Black T technique Aliquots: A 113.5 ppm B 113.5 ppm C 110.0 ppm The following treatments were carried out on the other three aliquots. l. Evaporation to dryness 2. Addition of 10 ml. of concentrated HN03 3. Evaporation to dryness 4. Solution with dilute HCl 5. Analysis for total calcium + magnesium concentration by above EDTA method Aliquots: A' 168 ppm B' 160 ppm C' 166 ppm Difference of x' - x = Z of complexed calcium and magnesium Thus: A' - A = 54.5 32% complexed 46.5 29% complexed Us I 0:! II 34% complexed C' - C = 56.0 .00 .75 .50 .25 .00 .75 .50 .25 .00 .75 .50 .25 .00 .75 .50 .25 .00 .75 .50 .25 .00 .75 .50 .25 .00 mm Calcium Ion Activity in Parts Per Million APPENDIX B 88 Valid for 3 x 10'2 M ionic strength solutions 1 x 10"3 M CaC12 = 0.00 mv. = 36 ppm +8. +8. +8 +8. +7. +7 +7. +7 +6. +6 +6 +6. +5 +5 +5 +5. +4. +4 +3. +3 +3. +3 +2. 75 50 .25 00 75 .50 25 .00 75 .50 .25 00 .75 .50 .25 00 .75 50 .25 .00 75 .50 25 .00 75 PPm 79 77. 75. 74 72. 71 69. 68 66. 65 63. 62 60. 59 58 56. 55 54 53 52 51 49. 48 47 46 .50 .25 .00 .75 .50 .25 .00 .75 .50 .25 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 ppm 45 44 43. 42. 41. 41 40 39 38 37 36 35 34 33. 33 32 31 30. 3O 29. 29 28. 28 27 26. U'IU‘I .75 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 .75 .00 .25 .50 .75 PPm 26 25. 25 24. 24 23. 23 22. 22 21. 21 20. 20 19. 19. 18. 18. 18. 17 17. 17 17 16. 16 16. 75 25 .75 .25 75 .50 25 Variations in calcium activities upon additions of other Experiment #1 : in grams 0.00 0.01 0.01 BaClz in grams 0.00 0.01 0.01 CuCl2 in grams 0.00 0.01 0.02 0.01 APPENDIX C 89 Ion Exchange Expe rime nts Rec onnai s sance in ppm 57 58 59 Ca2+ in ppm 56 54 53 Ca2+ in ppm 57 59 63 63 in grams 0.00 0.01 0.02 FeC13 ‘ tzO in grams 0.00 0.01 0.01 NiClZ in grams 0.00 0.01 0.01 0.01 cations Ca2+ in ppm 58 59 68 Ca2+ in ppm 57 57 57 Experiment #2: MgClZ added in grams OOOOOO Experiment #3: FBCIZ ' .00 .01 .01 .01 .01 .02 4HZO in grams 000000 .00 .01 .01 .01 .01 .01 Divalent Magne sium Ions Divalent Iron ppt. ppt. ppt.. Experiment #4: Divalent Barium BaClZ ' ZHZO in grams 000600 .00 .01 .01 .01 .01 .01 ppt. 90 in ppm 52 59 52 57 68 74 in ppm 56. 56. 56. 68. 68. 74. OOOWU‘IU'I A2 + in ppm 62 59 57 57 57 61 in ppm 52 57 62 62 65 74 in ppm 56. 62. 60. 68. 68. 68. OOOU‘IOU“! B (332+ in ppm 62 57 55 57 57 58 in ppm 52 55 55 62 62 74 in ppm 56. 60. 58. 68. 68. 68. 0000mm Experiment #5: FeCl3 - 6HZO in grams .00 .01 .01 .01 .01 .01 000000 Experiment #6: in grams .00 .01 .01 .01 .01 .01 000000 Trivalent Iron ppt. ppt ppt. ppt. Divalent Nickel * Ionic strength effect Experiment 7: in grams 0.00 0.03 0.03 0.03 Calcium electrode st Divalent Copper ppt. 91 Ca2+ in ppm 56. 56. 56. 56. 52. 47. OOU‘IUTU'IU'I A Ca2+ in ppm 52. 0 60. 5 68. 0 52. 0 Ca2+ in ppm 68 70 70 74 74 74 B Ca2+ in ppm 56. 55. 55. 59. 52. 47. OOOOOW B Ca2+ in ppm 52. 0 63. 5 68. 0 54. 0 Ca2+ in ppm 68 70 68 74 74 74 Ca2+ in ppm 52. 0 62. 0 68. 0 52. 0 arted a strong negative drift with precipitate. Expe riment #1: qq-Jxlq-qulxlxl-J-JQ-leoooooooooo pH .38 .32 .24 .15 .07 .98 .90 .84 .78 .71 .67 .62 .56 .52 .49 .46 .44 .43 .41 .40 Aliquot A: HCl in m1. OOOOOOOOOOOOOOOOOOOO NNNNNNNNNNNNNNNNNNNO 92 APPENDIX D pH Perturbations Acid titration of two aliquots of Red Cedar River Aliquot B analyzed four hours later than A I—‘HHF—‘HHh—‘HHHHHHHHHI—‘I—‘HH 100 m1. (332+ Activity .62x .62x .63x .66x .68x .70x .70x .71x .72x .72x .74x .76x .78x .78x .79x .79x .80x .80x .80x .80x 10'3 10‘3 10'3 10'3 10-3 10-3 10-3 10-3 10-3 10-3 10-3 10'3 10-3 10-3 10‘3 10'3 10'3 10'3 10-3 10'3 21312131313131313131ZIZIZIZIZIEIZIZIZIZI qqqqqqquqqqqqqqqmmmmmm pH .45 .34 .27 .18 .09 .01 .92 .85 .78 .76 .69 .65 .61 .58 .54 .52 .49 .47 .44 .46 .44 .43 Aliquot B: HCl in m1. OOOOOOOOOOOOOOOOOOOOOO NNNNNNNNNNNNNNNNNNNNNO WHHHHHHHHHI—‘I—‘HHI—‘I—‘HHHHHH 100 ml. Ca2+ Activity .62x .64x .66x .68x .68x .69x .71x .72x .73x .75x .76x .78x .78x .79x .80x .80x .80x .80x .81x .81x .82x .82x 10‘3 10-3 10-3 10-3 10"3 10‘3 10-3 10-3 10-3 10-3 10-3 10'3 10'3 10-3 10"3 10'3 10-3 10'3 10-3 10"3 10'3 10-3 K ZIZIZIZIZIZIZIZIZIZIZ:213121233123EIEIZIZ 93 Experiment #ZA: Acid titration of 75 ml. Red Cedar River sample with 2 x 10-2 M HCl pH HCl Ca2+ pH HCl Ca2+ added in ppm added in ppm in ml. in ml. 8.36 0.0 75.5 7.23 0.1 97.5 8.36 0.1 75.5 7.21 0.1 97.5 8. 26 0.1 77. 5 7.16 0.1 97. 5 8.17 0.1 77.5 7.20 0.1 97.5 8.08 0.1 79.0 7.20 0.1 97.5 8.04 0.1 81.0 7.14 0.1 97.5 7.94 0.1 81.0 7.19 0.1 97.5 7.86 0.1 81.0 7.11 0.2 97.5 7.81 0.1 81.0 7.10 0.2 97.5 7.80 0.1 81.0 7.04 0.2 97.5 7.76 0.1 83.0 7.01 0.2 97.5 7.73 0.1 85.0 6.96 0.2 97.5 7.66 0.1 85.0 6.98 0.2 97.5 7.62 0.1 89.0 6.90 0.2 97.5 7.58 0.1 89.0 6.90 0.2 97.5 7.58 0.1 89.0 6.88 0.2 97.5 7.55 0.1 89.0 6.85 0.2 97.5 7.50 0.1 89.0 6.77 0.2 97.5 7.49 0.1 89.0 6.87 0.2 97.5 7.50 0.1 91.0 6.75 0.2 97.5 7.44 0.1 91.0 6.70 0.2 97.5 7.46 0.1 91.0 6.68 0.2 97.5 7.45 0.1 91.0 6.58 0.2 97.5 7.39 0.1 93.0 6.58 0.2 97.5 7.35 0.1 93.0 6.57 0.2 97.5 7.38 0.1 91.0 6.55 0.2 97.5 7.38 0.1 91.0 6.52 0.2 97.5 7.34 0.1 93.0 6.46 0.2 97.5 7.34 0.1 93.0 6.48 0.2 97.5 7.33 0.1 93.0 6.44 0.2 97.5 7.32 0.1 93.0 6.37 0.2 97.5 7.33 0.1 97.5 6.32 0.2 97.5 7.27 0.1 97.5 6.26 0.2 97.5 7.28 0.1 95.0 6.18 0.2 97.5 7.32 0.1 97.5 6.11 0.2 97.5 7.23 0.1 97.5 6.04 0.2 97.5 7.27 0.1 97.5 5.98 0.2 97.5 94 Experiment #ZB: Base titration of 75 ml. aliquot of Red Cedar River sample, with 2 x 10’2 M KOH (continuation of #2A) pH KOH Ca2+ in ml. in PPm 8. 38 0.0 74.0 8.35 0.1 74.0 8.42 0.1 74.0 8.45 0.1 72.5 8.49 0.1 71.0 8. 53 0.1 68.0 8.61 0.1 68.0 8.62 0.1 68.0 8.67 0.1 65.0 8.73 0.1 63.5 8. 75 0.1 62.0 8.80 0.1 62.0 8.83 0.1 59.0 8.86 0.1 59.0 8.88 0.1 56.5 8.90 0.1 56.5 8.95 0.1 56.5 8.86 0.1 56.5 8.98 0.1 54.0 9.00 0.1 54.0 ppt. now clearly visible 9.03 0.1 53.0 9.06 0.1 53.0 9.06 0.1 52.0 9.08 0.1 52.0 9.11 0.1 52.0 9.13 0.1 49. 5 9.14 0.1 48.0 9.16 0.1 48.0 9.18 0.1 47.0 9.20 0.1 47.0 9.21 0.1 47.0 9.23 0.1 46.0 9. 24 0.1 45.0 Drop in calcium activity due in part to dilution, precipitation. \0\O\O\O\O\0\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\D\O pH .25 .27 .28 .30 .31 .34 .34 .35 .37 .39 .39 .41 .42 .44 .45 .47 .47 .49 .50 .52 .53 .54 .56 .57 .58 .59 .61 .61 .62 .64 .65 .66 .68 .68 KOH in m1. OOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOO y—ap—JHHHp—ap—IHp—IHHHHHp—aHHHHHHHHHHp—IHHHp—AHHHp—o Ca2+ in PPm 45. 43. 43. 43. 43. 42. 42. 42. 41. 41. 41. 40. 40. 40. 40. 40. 38. 38. 38. 38. 36. 36. 36. 36. 36. 36. 34. 33. 33. 33. 33. 33. 33. 31. OOOOOOOOOOOOOOOOOOOOOOOOU‘lmmU‘lmU'IU‘IU'IU‘IO \O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\O\D\O\O\O\O\O\O\D\O\O\O\D pH .69 .70 .72 .73 .74 .75 .76 .77 .78 .79 .82 .82 .82 .83 .84 .86 .87 .88 .89 .90 .91 .92 .93 .94 .95 .98 .99 .99 .99 .02 .03 .03 .03 .05 KOH in ml. and later to OOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOO HI—‘h—‘h—‘b—‘I—dt—‘i—‘D—it—‘D—‘r—‘I—‘HHb—‘I—iI—‘D—‘HHHI—‘i—‘fi-‘F—‘I—‘HP—‘r—lt—HD—‘I—iI—l Ca2+ in PPm 31. 31. 31. 31. 30. 30. 30. 30. 30. 29. 29. 29. 29. 28. 28. 28. 28. 28. 28. 28. 28. 26. 26. 26. 26. 25. 25. 25. 25. 25. 25. 24. 24. 23. U'IU‘IUTOU1U1U1U1U‘IU1U'IU1UTOOOOOOU'IUTOOOUTOOOOOOOCO THE" 95 Experiment #3: Titration of a Red Cedar River 100 m1. sample using 2 x10‘2 MHCl or KOH pH KOH Ca2+ added activity in ml. 8.38 0.0 1.78x10'3M 8.43 0.2 1.75x10'3M 8.53 0.2 1.72x10'3M 8.57 0.2 1.71x10'3M 8.63 0.2 1.70x10-3M 8.68 0.2 1.68 x10-3 M 8.73 0.2 1.68x10-3M 8.77 0.2 1.63x10-3M 8,32 0.2 1.62x10‘3M (3.36 0.2 1.60x10'3M 8,90 0,2 1.60x10’3M 3,94 0.2 1.60x10'3M 8.97 0.2 1.56 x10'3 M 9,00 0,5 1.55x10-3M 9,07 0,5 1.50x10'3M 9.20 0.5 1.48X10-3M 9.26 0.5 1.43x10'3M 9.31 0.5 1.41x10'3M 935 0,5 1.38::10'3 M 9.39 0.5 1.35x10-3M 9.44 0,55 1.32x10’iM 9.49 0.5 1.30x10' M Experiment #3: (cont'd.) \lxlxlxlflxlxlxlxlxlxlxlxl-xlxlxlxlxlxlxlxl\I-flmmmmmmmmmm pH .38 .36 .33 .30 .25 .21 .16 .11 .07 .02 .98 .93 .89 .86 .81 .77 .74 .70 .67 .64 .58 .54 .49 .43 .40 .35 .33 .29 .25 .23 .20 .18 .15 COOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOO IiCl added inInL NNNNNNNNNNNNNHI—nr—II—I—II—II—II—II-aI—II—Ir—II—II—II—II-dI-dI—iI—ao I—‘I—‘t—‘I—JI—‘I—‘HHI—IHI—‘HHHF—‘HHHHI—‘HI—‘D—‘HI—‘r—‘HF—‘HHHHH Ca2+ activity .78:x .75:x .78:x .77): .78): .79): .80:x .79); .79:x .79:x .85:x .82.x .83 .84 .83 .82 .83 .85 .85 .85 .85 .85 .85 .85 .88:x .88): .88:x .89}: .89): .89:x .89:x .89}: .89:x NRXNNXXNXNR X 10‘3 10-3 10-3 10-3 10-3 10-3 10'3 10-3 10-3 10'3 10-3 10-3 10‘3 10-3 10-3 10-3 10-3 10-3 10-3 10-3 10'3 10'3 10-3 10-3 10-3 10-3 10"3 10-3 10-3 10-3 10-3 10-3 10-3 31213131313131213131313131213121313131313121313131313131313131313 C‘O‘O‘O‘O‘OO‘O‘OO‘O‘O‘C‘O‘O‘O‘O‘O‘O‘O‘O‘C‘O‘O‘C‘xl-\1\l\1\1\1\1 pH .12 .11 .08 .07 .05 .03 .01 .99 .98 .96 .94 .93 .92 .92 .90 .88 .87 .85 .85 .83 .82 .81 .80 .79 .76 .71 .67 .62 .58 .55 .50 .46 HOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOO liCl added hirnL OU'IU'ImU'IU‘IUTU'INNNNNNNNNNNNNNNNNNNNNNNN HHHI—‘Ht—‘F—‘HHHHI—‘h—‘HD—‘HI—‘HHF—Ih—‘r—‘P—‘F—‘Hr—‘r—‘r—‘Hb—‘HH Ca2+ activity .893: .90:x .90:x .90}: .90): .90){ .90x .90): .90): .90}; .90): .90x .90x .90): .91x .91x .92x .92x .93x .93x .93x .91x .92x .93x .93x .93x .93x .93x .93x .93x .93x .93x 10-3 10-3 10-3 10"3 10-3 10-3 10’3 10-3 10-3 10-3 10-3 10-3 10-3 10-3 10-3 10-3 10‘3 10‘3 10'3 10'3 10': 10 10-3 10-3 10'3 10‘3 10'3 10-3 10'3 10-3 10-3 10'3 Z1313131313131212131313131313131313131313131313 Z 313131313 §.§ ‘1» in: ma vuafltw 97 Experiment #4: Reversible titration of Red Cedar River, 100 m1. sample, with 2 x 10-2 M HCl and 2 x 10-7- M KOH pH HCl Divalent pH KOH Divalent added ppnn added ppn1 hirnL hindL 8.38 0.0 130 6.49 0.5 153 8.32 0.2 133 6.56 0.5 153 8.24 0.2 133 6.65 0.5 153 8.14 0.2 130 6.73 0.5 153 8.06 0.2 137 6.83 0.5 153 7.96 0.2 135 6.91 0.5 154 7.87 0.2 135 7.04 0.5 154 7.79 0.2 138 7.16 0.5 154 7.72 0.2 140 7.35 0.5 154 7.65 0.2 141 7.58 0.5 155 7.60 0.2 143 7.87 0.5 155 7.53 0.2 141 8.11 0.5 155 7.49 0.2 138 8.36 0.5 155 7.44 0.2 139 8.56 0.5 155 7.40 0.2 140 8.71 0.5 155 7.36 0.2 140 8.83 0.5 155 7.27 0.5 141 8.94 0.5 155 7.20 0.5 142 9.01 0.5 155 7.11 0.5 140 9.09 0.5 155 7.06 0.5 142 9.14 0.5 153 6.99 0.5 140 9.19 0.5 153 6.95 0.5 142 9.24 0.5 151 6.93 0.5 143 9.29 0.5 149 6.90 0.5 146 9.33 0.5 149 6.84 0.5 146 9.37 0.5 147 6.79 0.5 146 6.75 0.5 146 6.70 0.5 148 6.66 0.5 148 6.61 0.5 151 6.58 0.5 151 6.55 0.5 153 6.50 0.5 152 6.47 0.5 153 6.43 0.5 153 6. 39 O. 5 153 1.33%” ."1 .—. .-.._ _‘,.. 98 Experiment #5: Reversible titration of three 100 ml. aliquots of Red Cedar River sample with 2 x 10"2 M HCl and 2 x 10'2 M KOH Additions A B C in ml. pH Ca2+ pH Ca2+ pH Ca2+ in PPm in Ppm in ppm HCl 0.0 8.27 89.0 8.29 89.0 8.29 89.0 0.5 8.07 89.0 8.11 89.0 8.03 89.0 0.5 7.82 93.0 7.90 93.0 7.85 91.0 0. 5 7.70 97. 5 7.70 95.0 7.68 97. 5 0. 5 7. 54 97. 5 7. 56 97. 5 7. 56 97. 5 0. 5 7. 45 97. 5 7. 46 97. 5 7. 46 97. 5 0.5 7.34 97.5 7.35 97. 5 7.35 97.5 1.0 7.19 97.5 7.19 97.5 7.17 97.5 1.0 7.06 <97.5 7.08 97.5 7.05 97.5 2.0 6.83 97.5 6.85 97.5 6.84 97.5 2.0 6.66 97.5 6.68 97.5 6.68 99.5 2.0 6.53 97.5 6.53 97.5 6.54 97.5 KOH 1.0 6.71 97.5 6.72 97.5 6.71 97.5 1.0 6.89 97. 5 6.89 97.5 6.89 97.5 1.0 7.08 95.0 7.07 95.0 7.11 97.5 2.0 7.93 93.0 7.91 93.0 8.06 91.0 2.0 8.75 81.0* 8.82 81.0* 8.87 81.0* H 1 2.C0 7.87 89 0 7.94 89.0 8.04 89.0 2.0 7.21 91.0 7.24 91.0 7.25 91.0 2.0 6 88 93 0 6 90 93 0 6 88 93 0 >kprecipitate 99 Experiment #6: Reversible titration of three 50 m1. aliquots of Red Cedar River sample with 2 x 10'2 M HCl and 2 x 10-2 M KOH Additions A B in ml. pH Ca‘?‘+ pH Ca2+ pH Ca2+ in ppm in ppm in ppm HCl 0.0 8.29 83.5 8.31 83.5 8.32 81.5 0.5 7.96 83.5 7.91 85.5 7.93 85.5 0.25 7.81 87.5 7.77 89.0 7.78 87.5 0.25 7.64 89.0 7.63 89.0 7.65 89.0 0.25 7.48 89.0 7.42 89.0 7.50 91.5 KOH 0.25 7.79 89.0 7.70 89.0 7.61 89.0 0.25 8.12 89.0 8.01 89.0 7.93 89.0 0.25 8.41 89.0 8.32 89.0 8.25 89.0 HCl 0.25 8.18 9.15 8.10 91.5 8.04 89.0 0.25 7.96 91.5 94.0 7.81 96.0 0.25 7.75 97.5 7.69 96.0 7.63 97.5 0.25 7.58 97.5 7.56 97.5 7.52 97.5 100 Exgeriment #7: Reversible titration of 100 ml. aliquot of synthetic Red Cedar River with 2 x 10"2 M HCl and 2 x 10'2 M KOH pH Additions Ca2+ in ml. in ppm HCl 8.08 0.00 97.5 4. 77 0. 50 pH too low for reading KOH 5. 54 0. 05 pH too low for reading 5. 85 0. 05 pH too low for reading 6.44 0. 05 89.0 6.70 0.05 97.5 7.05 0.05 99.5 7.17 0. 05 99. 5 8.10 0.10 99. 5 8.66 0.10 99.5 9.01 0.10 97.5 HCl 8. 50 0.10 97. 5 7.21 0.10 99.5 7. 03 0.10 95.0 6.58 0.20 97.5 KOH 7.71 0.30 95.0 101 Experiment #8: Reversible titration of 75 ml. sample of synthetic Red Cedar River using 2 x 10-2 M HCl and 2 x 10'2 M KOH pH Additions Cd2+ in ml. in ppm iic1 7.86 0.00 97.5 7.20 0.05 97.5 6.45 0.05 99.5 5.78 0.05 99.5 Kcui 6.03 0.05 99.5 6.38 0.05 99.5 6.66 0.05 99.5 6. 88 0, 05 99. 5 7.35 0.05 99.5 7.96 0.05 99.5 8.50 0.05 99.5 8.87 0.05 99.5 9.12 0.05 99'5 9.37 0.10 97-5 9. 54 O. 10 97. 5 9.62 0.10 97-5 9.84 0.20 97.5 9.96 0.20 97-5 10.06 0-20 97'5 10.12 0-20 97'5 10.45 0.50 95-0 PP“ 10.55 0.50 91.0 ppt Time 1/2 hour No H202 2 hours NO H202 6 hours 24 hours 2% H202 48 hours 102 APPENDIX E Eh - pH Perturbation Experiment Treatments with 30% H202 and 2 x 10‘2 M HCl 50 ml. aliquots of Red Cedar River Three replicates per block Calcium Activity Results No HCl 2% HCl/vol. 5% HCl/vol. 65, 65, 65 65, 65, 65 68, 67, 67 65, 65, 65 64, 64, 64 74, 79, 78 64, 64, 64 66, 66, 66 67, 67, 67 68, 68, 68 71, 69, 71 78, 78, 79 64, 64, 64 69, 69, 69 70, 70, 68 68, 68, 68 71, 71, 72 79, 79, 79 71, 70, 71 73, 73, 73 73, 74, 74 75, 75, 75 78, 78, 78 80, 80, 82 77, 77, 77, 78, 78, 78 84, 86, 84 88, 89, 89 94, 94, 96 94, 94, 94 10% 71, 83, 71, 84, 77, 87, 85, 95, 95. 99, HCl/Vol. 71, 71 83, 83 72 84 71, 84, 77 87 77, 87, 85 93 85, 93. 93. 95 99.101 103 Hum chem: co .w .00 .w do .w .mo .w om.m .38 £48 .648 NH .N. .Na .N. 00 .N. .mo .N. 8m 88m .mod “Ho ow nww ow .N¢.w Gm gm mfidmom mm H03 oz .3: Nomm oz .3: Nomm gem mydOfi wv .Ho>\ NONE oZ .BZNONm .3 mhsofi «N .331on oz .Ho>\ NONE e\eN mudos b .39/1on oz .3: Nomm .3 9353 N .3: Nomm oz 49,}on .NeN 32 N: 05:. 085 4016 m m3 "0 1111 3 mg "2 31 IIHHIIHIIH