ROBERT A. TAFT WATER RESEARCH CENTER REPORT NO. TWRC-1 DILUTE SOLUTION REACTIONS OF THE NITRATE ION AS APPLIED TO WATER RECLAMATION ADVANCED WASTE TREATMENT LABORATORY-I U.S. DEPARTMENT OF THE INTER/OR FEDERAL WATER POLLUTION CONTROL ADMINISTRATION OHIO BASIN REG/ON CINCINNATI, OHIO ------- DILUTE SOLUTION REACTIONS OF THE NITRATE ION AS / APPLIED TO WATER RECLAMATION by Frank C. Gunderloy, Jr., Cliff Y. Fujikawa, V. H. Dayan and S. Gird for The Advanced Waste Treatment Research Laboratory Robert A. Taft Water Research Center This report is submitted in fulfillment of Contract No. 14-12-52 between the Federal Water Pollution Control Ad- ministration and Rocketdyne, a Division of North American Rockwell Corporation. U. S. Department of the Interior Federal Water Pollution Control Administration Cincinnati, Ohio October, 1968 ------- FOREWORD In its assigned function as the Nation's principal natural resources agency, the United States Department of the Interior bears a special obligation to ensure that our expendable re- sources are conserved, that renewable resources are managed to produce optimum yields, and that all resources contribute their full measure to the progress, prosperity, and security of America — now and in the future. This series of reports has been established to present the results of intramural and contract research carried out under the guidance of the technical staff of the FWPCA Robert A. Taft Water Research Center for the purpose of developing new or im- proved wastewater treatment methods. Included is work conducted under cooperative and contractual agreements with Federal, state, and local agencies, research institutions, and industrial organi- zations. The reports are published essentially as submitted by the investigators. The ideas and conclusions presented are, therefore, those of the investigators and not necessarily those of the FWPCA. Reports in this series will be distributed as supplies per- mit. Requests should be sent to the Office of Information, Ohio Basin Region, Federal Water Pollution Control Administration, 4676 Columbia Parkway, Cincinnati, Ohio 45226. 11 ------- , ACKNOWLEDGEMENT This report is submitted to the Department of the Interior, Federal Water Pollution Control Administration, in fulfillment of Contract lU-12-52, Article III, B. The work was carried out over the period 5 January 1968 through 5 August 1968 by the Chemical and Material Sciences Branch of Rocketdyne Research Division. Dr. B. L. Tuffly (Manager, Environmental Sciences and Technology) served as the Program Manager. Dr. F. C. Gunderloy (Principal Scientist, Inorganic and Organometallic Chemistry) was the Responsible Scientist. Members of the Technical Staff contributing to this program were Dr. C. Y. Fujikawa, Dr. V. H. Dayan, Dr. S. R. Webb, Dr. M. A. Rommel, Dr. J. Foster, Mr. S. Cohz, and Mr. S. Gird. Dr. R. B. Dean (Ultimate Disposal) of the Cincinnati Water Research Laboratory acted as Project Officer for the Federal Water Pollution Control Administration. The Rocketdyne staff wishes to express their appreciation for the interest, expertise, and guidance provided by Dr. Dean throughout the course of this research. Thanks are also due to Mr. F. M. Middleton and Dr. D. G. Stephan of the FWPCA for time devoted to several enlightening discussions held at the Cincinnati Water Laboratory prior to the inception of this program. 111 ------- CONTENTS FOREWORD ii ACKNOWLEDGEMENT iii CONTENTS iv ABSTRACT vi INTRODUCTION 1 SUMMARY AND CONCLUSIONS 3 LITERATURE SURVEY 6 DISCUSSION AND RESULTS 3 BACKGROUND 3 Nitrates in Water - Occurrence and Effects 8 Natural Denitrification 9 Nitrate Removal Processes 9 SURVEY OF CANDIDATE AGENTS 10 Chemicals in Water Reclamation 10 Choice of Reducing Agents 10 Choice of Deammination Agents 13 Catalysis in Reduction and Deammination 1^ DILUTE SOLUTION REACTION STUDIES lU Screening of Reducing Agents 1^ Deammination Agents 16 iv ------- FERROUS ION AS A DENITRIFICATION AGENT ............ 18 Initial Screening Studies ........... jo Attempted Identification of "Missing N" .......... 20 Effect of pH on Denitrification .............. 22 Effect of Catalysts on Denitrification .......... 2h Effect of Air on Reduction and Denitrification ...... 24 Lime for pH Adjustment ................ \ 26 Effect of Phosphate and Carbonate Ions .......... 26 ANALYTICAL SUPPORT ....................... 3Q Background and Selection of Methods ............ 30 Analytical Program ................. 31 Development of Ultraviolet Method ............. 31 Ammonia Distillation Method ................ 37 EXPERIMENTAL DETAILS .................... APPARATUS ........................ ,g REAGENTS ....................... . PROCEDURE TEST RESULTS REFERENCES 45 APPENDIX - ANALYTICAL METHOD a. ------- ABSTRACT A new and unexpected partial denitrification of dilute nitrate ion solutions (10 to 50 ppm NO^-N) has been achieved by treatment with 8 moles of ferrous sulfate per mole of nitrate in unbuffered alkaline reactions. The nitrogen loss, which probably represents evolution of N2 or N20, has been as high as 50$. Total reduction to lost nitrogen plus nitrite and/or ammonia has approached 100$. The reduction takes place in the presence of partially oxidized black iron hydroxides, and requires catalytic quantities of cupric ion. Denitrification is suppressed by phosphates, as well as by several other factors, some as yet unidentified. Silver ion catalysis or a 16-fold excess of the ferrous salt permits reduction to ammonia in the presence of phosphate, but there is no accompanying denitrification. Keywords : Nitrates, wastewater, denitrification, reduction, ferrous ion, catalysis, cupric ion VI ------- INTRODUCTION The presence of nitrate ion in water, reclaimed or otherwise, presents several distinct problems. In high concentrations, it can cause methemo- globinemia, a disease of the newborn, and it serves as a nutrient for algae at any concentration. Algae bloom becomes a problem in many aspects of water usage, such as the fouling of reclaimed waters in reservoir storage. Although there have been numerous studies of methods of nitrate removal and control for use in water reclamation, removal by some direct chemical reaction, other than ion exchange, does not seem to have been given any really serious consideration in the past. This is not an oversight on the part of the interested scientific community; it is simply a reflection of circumstances that are not obviously amenable to attack by way of some common chemical reaction. For instance, if one considers a nitrate-N level of 10 ppm in water, then such ordinary reactions as quantitative precipita- tion or nitration of an organic compound cannot be brought to bear on removing nitrate from reclaimed waters. Precipitation requires expensive organic precipitants (e.g., "Nitron"), and nitration of organics is only accomplished in concentrated solutions. Conversion of nitrate to a nitrogenous gas and denitrification by the evolution of such a gas is an attractive concept, and there are three gases that can be considered, although each has drawbacks (Ref. A1-A3). These are ammonia, nitrous oxide (N20) and nitrogen. (Such gases as N02, NF3, NOC1, NO, and the like have to be discarded on the basis of reactivity, water solubility, or the simple impossibility of converting N0§ to such a gas in any simple reaction or series of reactions.) Ammonia has high water solubility, and requires a physical stripping step. It cannot simply be left in the water, since eventually, as part of the biological nitrogen cycle, it will be reconverted to nitrate. Nitrous oxide, while not nearly as soluble as Nfo, still has a fairly high water solubility compared to nitrogen. Nitrous oxide in solution does oxidize, albeit at a very lov; rate (Ref. A-l). Thus, at concentrations in the ppm range, N20 could be subject to the same drawbacks as ammonia. As far as chemical characteristics are concerned, nitrogen itself would be the ideal choice, since it has the lowest water solubility and greatest oxidation resistance of all the nitrogenous gases. However, in the oxynitrogen and hydronitrogen series of compounds, nitrogen is unique in that it is very difficult to obtain by simple reduction (Ref. A-3).." While redox potentials often appear favorable, in the case of N0§ there is a large activation energy that must be overcome, and very strong reducing agents lead to NHo in most cases, while weaker ones take the nitrogen only to the +3 oxidation state. Nitrogen evolution would be an attractive means of ridding water of nitrate ion, and the work described in this report was undertaken with ------- Just that objective in mind, even though there appeared to be no simple one-step reaction sequence to go from N0§ to N2. The premise ofVthis work was that there might be efficient and economically attractive two- step reaction sequences that could be applied, based on reduction of nitrate to nitrite ion, and subsequent deammination of a primary amine with the nitrite. A model reaction series for this concept, using formal- dehyde as the reducing agent and urea as the deammination agent is shown in equations (l) and (2) . + HCHO •* 2HN02 2HN02 + COdlH,,),, •» 2N2 * C02 + 3^0 (2) If applicable to dilute solutions of nitrate ion, these known reactions would be very attractive. The products N2, C02, and IfeO are all innocuous, and based on bulk prices, a chemicals cost of 2 to 3 cents per thousand gallons could be projected for removal of NO§-N at the 10-20 ppm level. Further consideration of the model reaction sequence allows one to speculate about other potential advantages of a direct chemical reaction Incorporation of such a sequence into an existing plant might be possible by simply adding the appropriate metering devices to inject the reagents into the stream. A chemical process is easily adjusted to variable nitrate levels; there is no need to design for the maximum level, which would result in unused removal capacity during minimum flow periods. A chemical process can be turned on and off at will; the plant that experiences seasonal nitrification would be greatly benefited by having a process that it could start and stop on demand at the appropriate time of the year. Thus, flexi- bility and inherent reliability could make the reduction-deammination sequence more attractive than just simple economics might suggest. While many reduction and deammination reactions are known and have been extensively studied, the conditions for such studies have been almost entirely restricted to relatively concentrated solutions. Dilute solution reaction chemistry of any type, let alone dilute solution reaction chemistry of specifically nitrate reduction and deammination, is an area of study which has been faintly touched at best. The objective of the present program was to select groups of candidate reduction and deammination agents that might possibly be used in water reclamation processes, and to test the feasibility of developing a denitrification process using these agents when the nitrate ion was present at dilutions of 10 to 50 ppm of NO"-N. ------- SUMMARY AMD CONCLUSIONS The objective of this program was to demonstrate the feasibility of denitrification by a chemical process. Although departing somewhat from the original reduction-deammination concept, such a demonstration was achieved in the course of the work. After an extensive literature survey, eight reducing agents and three deammination agents were selected for testing the feasibility of denitrification by a reduction-deammination process. The original plan was to study various pairs of these agents in a statistically designed matrix of experiments, varying numerous environmental factors such as pH, temperature, etc. However, it soon became evident that unknown second- order interactions in the design could defeat the purposes of such a study. Accordingly, these eight reducing agents were screened under anaerobic conditions at high NO^-N concentrations (50 ppm) and high temperature (85° F) with the agent in excess, varying only pH and catalyst. On this basis, ferrous ion (Fe++), iron powder, and hydrazine or its salts (N2Hij, ^1^30^) showed appreciable reducing power. A very small amount of reduction was accomplished with glucose. Formaldehyde, carbon, sulfur dioxide, and carbon monoxide were inactive. The deammination agents studied were sulfamic acid (HSC^NHg) and urea. Glycine was slated for study, but later abandoned. No nitrogen loss was detected that could be attributable to the deammination agents in any of the tests where reducing and deammination agents were studied together. Separate studies showed that urea was ineffective, and that sulfamate could only deamminate under acid conditions. Of the three reducing agents that passed the screening, only ferrous ion was economically attractive; it is available as crude copperas, FeSOli'TH^O, for only four dollars per ton. From the technical viewpoint, ferrous ion was the only choice, since the test series revealed that up to 55% denitrification was occurring with ferrous ion alone, apparently by direct reduction of NO" to either N2 or NgO ("missing N"). Accordingly, the bulk of the experimental program was devoted to studies of the direct denitrification reaction with ferrous ion. Ferrous sulfate was used as the source of ferrous ion, but the solution must be initially basic, so that the system is actually heterogeneous, with pre- cipitated ferrous hydroxide being the reducing agent. As the reaction progresses, the pH drops and the black ferrous-ferric complex is formed (Te^pii or the corresponding hydroxide). Assuming that the "missing N" evolves as Ng, then the reduction is a five-electron reaction, and since additional ferrous ion is consumed in forming the complex, the theoretical requirement for complete reduction would be 7.5 moles of Fe"*"1" per mole of NO" ------- At the 10 ppnHDvJI level, using an 8:1 molar Fe++:N<>; ratio, reduction (to mixtures of N02, Ah, and "missing N") ranged from 50)Tto looi and denitrification ranged from 1(# to UjJ. These results vere obtained over an initial pH range of 7 to 11 using either lime or sodium hydroxide for pH adjustment. Neither reduction nor denitrification was observed under acid conditions. A trend for the amount of reduction to increase with increasing pH was noted, but the amount of denitrification did not seem to follow a trend. Catalysis by either cupric ion (Cu++) or silver ion (Ag*) in 1-5 ppm concentration is necessary for reaction to occur, but denitrification was observed only when cupric ion was used as a catalyst. Addition of phosphate ion to the solutions interfered with catalysis by cupric ion, and no reduc- tion occurred, even when the catalyst level was increased. However at increased levels of iron (16:1 and 2k:l), or with silver as the catalyst, reduction occurred readily. As before, no denitrification occurred when Ag was the catalyst. Carbonate ion did not interfere with the reduction reaction, but, again, no denitrification occurred when C0| was present. Carbonate buf- fered the solutions; this fact, coupled with other results where pH was held relatively constant by addition of base during the course of a run indicated that denitrification would not occur unless the pH was dropping as the reactions progressed - that is, buffering prevented denitrification. Based on these results, development of a denitrification process based on direct chemical reduction of nitrate ion to nitrogen or N2© appears feasible. The denitrification can be carried out with an inexpensive reducing agent (copperas) and is worthy of further technical evaluation. However, much more will have to be known about the basic chemistry of this reaction before a definitive process emerges. The effect of pH and buffering on the reaction needs to be resolved since there is no evident explanation for the absence of denitrification under constant pH conditions. Quite possibly, observation of denitrifica- tion has been obscured by the effect of some unknown variable. The fact that the amount of denitrification does not correlate with pH level, and that there has been an occasional failure to denitrify even in the absence of buffering, indicates that such a variable may indeed exist. The hetero- geneity of the system could be the source of this inexplicable behavior. Minor variations in the mode of precipitation of the ferrous hydroxide,' the manner of absorption of catalyst ions and the degree to which this' precipitate absorbs them, the manner in which the ferrous-ferric complex forms, and a number of other factors relative to the solid phase could all play a role, and consistent denitrification and reduction might wen prove to be a function of consistent precipitation technique. ------- Of great importance, too, is a determination of the identity of "missing N", and an investigation of the intermediates that lead from HOo to "missing K". Obtaining total denitrification may veil depend on steps in the mech- anism that are not evident in the present vork. For instance, if N2 is the gas evolved, the denitrification could actually be the result of the oxida- tion of N^^h or NHgOH formed at some intermediate stage. The possibility of loss of nitrogen in some form as a part of the precipitated iron insol- ubles should also be investigated, even though such a mode of denitrifica- tion seems unlikely. Further studies are also required on catalysis of the reaction. Cupric and silver ions are the classical catalysts for homogeneous reductions, yet ferrous ion reduction/denitrification reaction is apparently heterogeneous. Cupric and silver ions are definitely different in their response to phos- phate, and apparently different in their ability to lead to denitrification, although this last difference could again be confounded vith some other variable. With the information now at hand, there are three possible approaches to development of a ferrous ion denitrification process as described below. The first approach, total denitrification with ferrous ion, assumes that further studies of the reaction will result in the data necessary to make reduction of HO^ to "missing N" by ferrous ion both consistent and close to Quantitative. In the second approach, a reduction-deammination sequence, reduction and some denitrification is carried out by the ferrous ion, and the pH falls from some_initially alkaline value to belov 7 as the reaction pro- gresses. If HOg vere the major reduction product other than WgO or Ng, then deammination might still be used to complete the denitrification. However, carbonate ion would have to be absent for this sequence to take place, which would severely limit the applicability of this process. The third possible process, coupling ferrous ion reduction with ammonia stripping, might be the easiest to develop. In some of the reactions conducted during this study, 33-^5$ denitrification occurred, and another 3^6$ of the nitrate was converted to ammonia. Thus, a sequential ferrous reduction-ammonia stripping process has demonstrated potential of up to 90$ total denitrification. ------- LITERATURE SURVEY Approximately one-third of the effort on this program was devoted to a comprehensive literature search, covering the period from early 1968 wen back into the nineteenth century. Recent references were obtained from Keywords (Chemical Abstracts), Chemical Titles. Current Contents. Water Pollution Abstracts, and on-the-shelf journals and reports. For thTperiod back through 1907, Chemical Abstracts and the annual literature reviews in tne Journal Water Pollution Control Federation were prime sources. Qnelins Handbuch der anorganischen Chemie revealed references back to the early - 1800's. A patent search was conducted by the North American Rockwell Patent Department, and several pertinent current references were provided by the FWPCA Project Officer, Dr. R. B. Dean. While the primary objective of the search was selection of suitable reduction and deammination agents for subsequent laboratory testing, much related information was collected on water reclamation in general, and nitrified waters in particular. Between kOQ and 500 references were col- lected in original or abstracted form. For this report, slightly more than 200 references to the most pertinent and informative articles are provided. The "Discussion and Results" section, which follows, contains ref- erences to the literature sources which, for the convenience of the reader are grouped as follows in the section entitled "References": ' A. Nitrates in Water General Chemistry 5-30 Occurrence and Effects 31-^8 Natural and Biological Denitrification U9-68 Elimination Methods B. Water Reclamation 1-12 Conventional and Tertiary Treatment 13-25 Iron Salts in Water Treatment 26-30 Carbon in Water Treatment C. Reducing Agents for Nitrate 1-23 Ferrous Salts 2lf-28 Carbon 29-33 Sulfur Dioxide 35-^7 Formaldehyde U3-56 Sugars 57-61 Powdered Iron 62-69 Hydrazine and its Salts 70-85 Miscellaneous ------- D. Deaiamination Agents 1-lU Sulfainic Acid 15-21 Urea 22-21* Amino Acids 25-3^ Miscellaneous E. Catalysis 1-6 (Wot subdivided) F. Analytical Methods 1-9 (Not subdivided) ------- DISCUSSION AND RESULTS BACKGROUND Nitrates in Water-Occurrence and Effects (See References A-5 through A-30) Nitrates occur in natural and reclaimed waters in amounts ranging from a fraction of a part_per million to several hundred ppm, calculated as nitrate-nitrogen (NO^-N). The sources may be natural, such as leaching from nitrate deposits (e.g., guano in limestone cave areas), the natural decay and oxidation of nitrogenous organic matter (protein) as carried out by certain microorganisms, and the fixation of atmospheric nitrogen as NO and NOg from electrical discharges during thunderstorms. The sources may also be man-made, such as leaching from agricultural lands treated with nitrogenous fertilizers, effluent from fertilizer manufacturing plants, and effluent from other chemical and manufacturing processes that employ nitrates in one form or another. In water reclamation, nitrate ion may find its way into the stream initially from any of the above sources. However, regardless of the initial water quality, secondary treatments based on biological oxidation of organic matter (i.e., activated sludge and trickling filter processes) can them- selves introduce additional nitrate. Hence, almost all secondary effluent contains an appreciable quantity of nitrate ion. There are two distinct problems associated with nitrates in water: methemoglobinemia and algae bloom. Methemoglobinemia is a serious and often fatal disease of the new-born, characterized by cyanosis, for which nitrates have been designated a causative factor. A drinking water standard of 10 ppm N03-N has been established by the Public Health Service (Ref. A-30) as a preventive measure for this disease. Both ammonium and nitrate ions, as well as phosphates, are excellent nutrients for plants, including algae (Ref. A-lU). Algae bloom is one aspect of a natural process called eutrophication, wherein a body of water such as a lake is gradually converted into a swamp and eventually a meadow. The beginning and end of this process are not particularly unpleasant; however, in the intermediate stages, algae and aquatic plants grow in abundance as nutrients build up, this abundant organic matter decays depleting oxygen and killing the aquatic fauna. Unfortunately, man-made sources of nutrients accelerate this process, and induce the algae bloom stage in receiving waters where it would not otherwise occur. Eutrophica- tion would be intolerable in a reservoir to be used for reclaimed water for drinking purposes. Other water uses, such as recreation, incorporate a certain esthetic value which is certainly not enhanced by overgrowths of algae. ------- Natural Denitrification (See References A-31 through A-1*8) Countering the natural and man-made nitrification processes are natural denitrification processes. The conditions for nitrification and denitrifi- cation are very similar, with the main difference being that denitrification occurs only when the water is close to being depleted of oxygen. Under such conditions, certain microorganisms will continue the oxidative decradation of organic matter, using the nitrate ion as the oxygen source, and eliminating the nitrogen as N2 gas. Denitrification can occur in water, in soils, and in accumulated masses of organic matter such as silage (Ref. A-37, A-^U). Good reviews of natural denitrification processes are presented by the Thames Survey Committee (Ref. A-l*7) and by Camp (Ref. A-3U). The latter author, however, implies that natural denitrification can be partly chemical, and uses the deaznmination reaction of nitrite ion with urea as an example, showing the reaction to be thermodynamically favorable. Chemical reduction by ferrous salts is also believed to play a role in natural denitrification in certain soils (Ref. A-31, C-22) and acid tropical waters (Ref. A-Uu). However, it is generally agreed that most natural denitrification is bio- logical rather than chemical (Ref. A-33). In the vater reclamation field, natural denitrification first came to attention as a problem (Ref. A-32, A-3b through A-l*2). The nitrogen evolved, when natural denitrification occurs in sedimentation basins, causes the phenomenon known as "rising sludge" or "rising hunus". That is, bubbles of nitrogen entrapped in the sludge carry it to the surface rather than allowing it to settle. In recent times, natural denitrification has been turned to good use, and forms the basis for several advanced waste treatment; processes (Ref. A-36, A-^3, A-l*6, A-55). These processes are reviewed below. Nitrate Removal Processes (See References A-l*9 through A-68) There have been a number of methods studied for producing denitrified water. These include: control of secondary treatment processes in such a manner that nitrification is minimised; ion exchange; extraction; and biological denitrification. Adsorption of nitrate on such substrates as carbon, alumina, and silica gel has also been noted, but does not seem to have been studied for the specific purpose of developing a removal process. An occasional excursion into reaction chemistry has been made (Ref. C-23, C-58) using ferrous salts and iron powder, but with little or no success. However, concentrates (primarily radioactive wastes) have been successfully treated with formaldehyde and sugars as reducing agents, as discussed later in this section. Controlling the secondary treatment processes can be effective, but suppression of nitrification is often accompanied by some undesirable effect, such as decreased BOD removal. (BOD, Biochemical Oxygen Demand, a measure of the biodegradable organic content of the water.) Ion exchange is very efficient, but costs can be high. Extraction works very nicely on ------- concentrates, but is inefficient for dilute nitrate solutions. Apparently none of these processes are under serious consideration at the present time for extensive incorporation into reclamation plants. A more detailed pic- ture of the state-of-the-art of these processes, as well as processes designed to remove ammonia, may be found in the recent publication of Farrell, Stern and Dean (Ref. A-55). Biological denitrification is currently under active study (Ref. A-U6) There are two modifications of this process. The first (Ref. A-36) involves a sequence which first nitrifies the water to the greatest extent possible then carries the water into an anaerobic chamber where sludge from a previous step is used to supply organic food to the denitrifying microorganisms. The second modification has grown from the unexpected denitrification observed when activated carbon columns were being studied for tertiary treatment (Ref. A-i*3). In this case, the denitrifying organisms had established themselves in the carbon columns. Supplying methanol as additional food for the bacteria increases the efficiency of this process, and sand has been successfully substituted for the carbon. SURVEY OF CANDIDATE REDUCING AMD DEAMMENATION AGENTS Chemicals in Water Reclamation (See References B-l through B-30) In choosing agents for the reduction-deammination study, attempts were made, wherever possible, to project the tentative process in terms of other current and future processes to see where the new process might be conveniently incorporated. In general, it appears that flocculation and coagulation will play an increasing role in the future, using either lime or alum. (See Ref B-l for a tabulation of well-developed tertiary treatments.) Lime flocculation can reduce phosphate ion concentration, so it is reasonable to expect the waters to be highly basic during some stage of reclamation processes aimed at controlling nutrient content. The lime is often used with another salt as an additional coagulant. Ferrous and ferric salts (Ref. B-13 through B-25) have been used in this manner. Finally, activated carbon may be used in the final stages to remove the last traces of organic matter (Ref. B-26 through B-30). Both ferrous salts and carbon are potentially reducing agents for N03 ion, and are discussed further below. Choice of Reducing Agents Eight reducing agents were selected for study as a result of the literature survey. The information amassed on each of these is summarized below. 10 ------- Ferrous Ion (Ref. C-l through C-23). Ferrous ion, as either ferrous sulfate or ferrous hydroxide, has formed the basis for many analytical determination of nitrate ion,Converting the N0§ to NH3. The first step, conversion to nitrite ion (N02), is slow but may be catalyzed by silver or cupric ions. Subsequent reductions proceed to ammonia by attacking N02 and NO from the HN02/N02/NO/H20 equilibrium. A summary of almost an possible reaction sequences is given in Ref. C-22. The reduction is most effective under faintly or strongly basic solutions, but can also occur under acid conditions. It has been stated that the concentration of ferrous ion must be at least 70 ppm in order to be effective (Ref. C-1+). Ferrous ion was once studied for use in water reclamation, and was shown to be capable of 90$ conversion of NO? to NHo at the 100 ppm NOo-N level (Ref. C-23). This work was abandoned because of the "ferrous and ferric hydroxide sludges" that formed. From an operational viewpoint, this is difficult to understand, because the "sludges" are heavy, settle easily, and so should be easy to separate from the treated water. However, ultimate disposal of the "sludges" may pose some problems (Ref. B-13), and will have to be given serious consideration in the development of any process based on ferrous ion. As noted earlier, some of the natural denitrification and reduction in water and soils has been attributed to the presence of ferrous salts (See page 9). Carbon (Ref. C-2U through C-20). Nitrate can be reduced to nitrite by carbon, with C02 being the other product. The well-known "wet-ashing" technique for removal of carbonaceous materials in analytical procedures is a good example. Coal, carbon black, and graphite reduce nitrate. How- ever, high concentration and heating are generally required, under which conditions mixtures of nitrogen oxides are evolved. Sulfur Dioxide (References C-29 through C-33). Sulfur dioxide, if shown to be suitable as a reducing agent, could play a dual role in water reclamation because of its bacteriostatic properties. Sulfur dioxide can reduce nitrate to various products, including hydroxylamine (Ref. C-33). The most pertinent article (Ref. C-31) claimed that a 10$ NH3 solution saturated with S02 until faintly ammoniacal would reduce oxidizing anions and completely eliminate nitrite ion as N2. Formaldehyde (References C-35 through C-U?). Formaldehyde has been extensively studied as a reducing agent for nitrate concentrates ("Purex wastes") by the Atomic Energy Commission. Under these conditions, NO 'and N02 are evolved, which implies that HNOg would be the product in dilute aqueous solution. There is an induction period in the reaction which can be overcome by ferric ion catalysis (Ref. C-U2, C-^3). 11 ------- One of the few articles dealing with dilute solution reduction of nitrate ion has some interesting information on formaldehyde (Ref. C-36). In the photochemical reduction of nitrate to nitrite, with the results V expressed as a ratio (R = — ?), it is shown that the R value is increased NOo by a factor of 3 to k by the addition of formaldehyde. Working with a 0.05$ KN(>3 solution, the value for R was 60 after 20 minutes of ultraviolet radiation in the presence of formaldehyde, representing a change in NO§-N concentration of about 70 ppm down to 1 to 2 ppm. The data did not explicitly show, however, that part of the reduction was by direct reaction with for- maldehyde rather than being totally photochemical in nature. Sugars (References C-U8 through C-56) . Sugars have also been used by the Atomic Energy Commission to reduce nitrate concentrates. Sugars also increase the R value in the photochemical reduction of NO? (see above). A sugar need not be one of the classical reducing sugars to react readily with nitrate: sucrose as well as glucose is suitable. Powdered Iron (References C-57 through C-6l) . Powdered iron is similar to ferrous ion in the manner in which it reduces nitrate. It was once examined as a reducing agent for eliminating nitrate ion from drinking water, but was classed as ineffective because it did not convert NOg directly to N2 (Ref. C-58). The efficacy of powdered iron as a reducing agent depends to some extent on its method of manufacture (Ref. C-59). Hydrazine and its Salts (References C-62 through C-69). Hydrazine is an excellent reducing agent for nitrates, and forms the basis for the Auto Analyzer now used for analysis of various waters (Ref. C-63, C-6?). The reader should note that this reduction reaction, if it is to be of value in the proposed process, must lead to near-quantitative oxidation of the hydrazine to Ng. If this does not occur, then the reaction may introduce more inorganic nitrogen compounds into the water than are removed by the overall process. Carbon Monoxide. Carbon monoxide was selected as an agent purely on the basis of economy and the desirability of having COg as a byproduct. No references to CO/N03 reactions were found; on the contrary, a very early reference (1851) clearly states that carbon monoxide and nitric acid do not react (Ref. C-83). Miscellaneous Reducing Agents (References C-70 through C-85). Other reducing agents noted during the course of the literature survey were generally metals and various lower- valent salts of metals (e.g., Al, Ti**, Cu-Cd, Sn"*"1", etc.). These were eliminated from consideration on the basis of undesirable byproducts and/or expense. 12 ------- Choice of Deammination Agents Three deanmination agents were selected, based on the information presented below. Sulfamic Acid (References D-l through D-lU). Based on an article com- paring sulfamic acid and urea (Ref. D-l^), sulfamic acid was the prime choice for deEmmination. This article showed quantitative reaction of the theoretical amount of sulfamic acid and nitrite ion in two minutes even at the 2 ppm NOjj-N level, whereas even excess urea could not effect complete deammination at that nitrite concentration after one hour. The system sulfamic acid-nitrite ion has been studied extensively in analytical appli- cations, in terns of the intermediates formed, and with respect to thermo- dynamics. A means for removing nitrite ion from boiler water using sodium sulfamate has been patented (Ref. D-13). It should be noted that sulfamic acid can react with nitrate ion as well as nitrite, yielding ^0. However, this reaction is very slow below 60° C (Ref. C-7). Urea (References D-15 through D-21). As noted above, urea is not nearly as effective as sulfamic acid for deammination, although its use has been patented for treating waters to remove nitrite ion (Ref. D-15, D-21). Amino Acids (References D-22 through D-2U). Since amino acids are present in water during various stages of reclamation processes, they might provide an in situ source of deammination agent. Glycine was chosen as the model originally, since it reacts readily with nitrate, although side reactions give some C02 and N2<> as well as nitrogen (Ref. D-22). Compara- tive studies (Ref. D-23) have shown that relative to alanine (1.00), most amino acids, glycine included, are deamminated at about the same rate (0.70 to 1.50), although there are extremes such as cystine (3-12) and isovaline ( Miscellaneous Deammination Agents (References D-25 through D-3^). In general, all primary organic amines and amides are susceptible to deammina- tion by NOjj. Azides also react with nitrite in a similar fashion. How- ever, none of these offer any particular advantage over the agents dis- cussed earlier, and, of course, the organic amines will leave an organic residue, which is not desirable. Ammonia and its salts generally require heating in order to react with nitrites at appreciable rates. However, it was tentatively planned to test ammonia in combination with 502, in view of the data presented earlier (p. 11) on this particular combination. 13 ------- Catalysis in Reduction and Deanmination (See References E-l through E-6) Catalysis plays a role in many of the reactions mentioned earlier, i.e., iron and ferrous ion reductions are catalyzed by cupric or silver ions, hydrazine reductions by cupric ion, formaldehyde reductions by ferric ion. Catalysis has a role in certain biological nitrate reductions as well (Ref. E-3). Cupric ion and silver ion are the classic homogeneous reduction catalysts, particularly for homogeneous reductions with hydrogen (Ref. E-U). Of interest to the present program is the fact that catalytic activity of these ions may be enhanced more than a hundredfold in the presence of organic acids. Magnesium, cadmium, and zinc ions are also reported to be active catalysts in the presence of organic acids. An article published in 1923 (Ref. E-5) reports that nitrate ion reduction is an autocatalytic process, and that in the absence of some small initial concentration of nitrite ion, nitrate cannot be reduced by ferrous ion, formaldehyde, mercurous ion, or a number of other agents. Nitrite ion was removed from the test solutions with urea and amino acids. If this phenomena were indeed confirmed, it would mean that reduction and deanmination would have to be sequential operations. However, the conclusion is refuted to some extent by other work; the simultaneous reduction-deammina- tion cited earlier for m^/SO^ solutions is one example. DILUTE SOLUTION REACTION STUDIES Screening of Reducing Agents The initial laboratory plan in this program called for a series of statistical test matrices, studying reducing-deammination pairs, with high and low levels of the variables as given in Table 1. Each test solution was to be sampled after reaction times of one hour and 2h hours. Each sample was to be analyzed for nitrate, and if reduction had occurred, for nitrite and ammonia to determine the total nitrogen balance. In selecting the levels of variables, it was assumed that the denitri- fication treatment would be applied to secondary effluent, and the values for pH and temperature are believed to be the extremes. A different pH range was chosen for ferrous ion, Fe^, since in this case it was assumed that ferrous treatment would be coupled with lime coagulation, and the denitrification would thus be carried out in alkaline waters. The initial tests within the first matrix, if accepted at face value would have led to the strange conclusion that ferrous ion could not even ' ------- TABLE 1 REDUCTION-DEAMMINATION REACTION VARIABLES Variable High Level Low Level NO"-W Concentration Temperature PH ^ (Fe excepted) (with Fe**) Reducing Agent Concentration Deanmination Agent Concentration Order of Addition Atmosphere Catalyst 50 ppm Ol-O T, 9 11 Threefold excess Threefold excess Deammination agent added one hr after reducing agent Nitrogen (water deaerated) 1 to 5 ppia of Cu++ or 10 ppm 6 8 Stoichionetric Stoichiometric Agents added simultaneously Air (water untreated) No catalyst 15 ------- reduce nitrate ion, let alone participate in a electrification sequence. In view of the large amount of literature that had been uncovered on the reducing power of ferrous salts, this result appeared to be anomalous. In statistical terms, the apparent anomaly arose because the results were confounded by a second-order interaction of variables. The matrix design was such that air oxidation of ferrous ion was predominant in cat- alyzed tests, and the rate of the uncatalyzed reactions was essentially nil at the low concentration levels. Since much less information was available on the other seven reducing agents, it was decided that statistically designed matrices with a large number of variables could not be conducted with any degree of assurance that other unrecognized interactions would not lead to fallacious conclu- sions. Accordingly, a new test series was devised to screen the'reducing agents with most of the variables fixed at the high levels. Only the effects of pH, catalyst, and deaamination were studied in this new series, with the deammination agent added either initially after 2k hours, or not at all. However, the deammination agents had no effect on these reactions, as discussed later in this section. The results of the new tests, which amounted to screening the candidate agents for reducing power under the most favorable conditions possible, are given in Table 2. On the basis of these tests, only ferrous ion (Pe**) appears to be a potentially useful agent. Hydrazine and its salts, since they did not achieve quantitative reduction, introduced more inorganic nitrogen than they could remove. Iron powder is considerably more expensive than Fe"*"*" ion. The latter is available as crude copperas, at a cost of a few dollars (less than $5.00) per ton, while the most optimistic estimates for iron powder are in the area of 10 cents per pound. Deammination Agents Deammination agents (urea, or sulfamic acid initially neutralized to avoid additional pH adjustments) were present in many of the tests cited above, but once Fe ion had emerged as the prime reducing candidate, the deammination phase of this effort was stopped. The Fe++ reduction must be carried out under initially basic conditions, and sulfamate deamminates effectively only under strongly acid conditions. This conclusion was reached during a separate series of studies carried out as part of the supporting analytical effort (See p.35). The poor performance of urea, as cited in the literature (Ref. D-lU), was also confirmed. At the 3 ppm NO|-N level, no deammination could be detected after U8 hours at either pH U.O or 10.6, even though the urea was present in threefold excess. Glycine was slated to be studied as a deammination agent, but was by- passed as the investigation of ferrous ion proceeded. 16 ------- TABLE 2 REDUCING AGENT SCREENING Reducing Agent Conditions Varied (and number of tests) Reduction Observed, % so2 Carbon CO Glucose Fe Powder Fe Ion Fe , Cu catalysts; pH 6 and 9 (k tests) _ _ . . Fe , Cu catalysts; pH 6 (U tests) Fe' ' , Cu catalysts; pH 6 (U tests) Cu catalyst; pH 6 (2 tests) _ +++ _ ++ ,,+5 . , . Fe , Cu , V catalysts; pH U and 6 (6 tests) Fe , Cu catalysts; pH 11 (U tests) j, j^ Cu catalyst, pH 11 (U tests) I i^ Cu catalyst, pH 6 (3 tests) *Cu catalyst, pH 11 (6 tests) None None None None 3 to 6 (with V+5 only) 10 to UO (none with Fe 10-20 15-^5 15-^0 +++ Fixed Conditions: Temperature: Test Time: Atmosphere: NOl-N Concentration Reducing Agent Concentration: 85° F U8 hr. Nitrogen 50 ppm 3X (Denotes 3 moles of reducing agent for each mole of NOZ) * Tests with Fe slightly more complex. Discussed in more detail in the text. 17 ------- FERROUS ION AS A DENITRIFICATION AGENT Initial Screening Studies As noted in Table 2, the initial screening tests with ferrous ion were slightly more complex than those with the other reducing agents. The initial matrix+tests had indicated that bubbling air through the solutions oxidized the Fe before it could reduce NOo in cases where catalyst was present, and that uncatalyzed reactions were not possible at high dilutions. However there was still the possibility that there was an interference from the deanmlnation agent, either direct or by suppression of the autocatalytic influence of H02 ions, as had been indicated in the literature (Ref E-5 See p. 1U for discussion). Accordingly, the first screening tests with Fe*+ incorporated the sulfamate addition time and presence or absence of N05 ion as variables. Test conditions and results are given in Table 3. The above data indicate that nitrite does not catalyze the reduction and that sulfamate does not interfere with the reduction. If anything sulfamate aids reduction; more reduction occurred in Tests 3 and U where sulfamate was added initially. Of course, under the basic conditions, sulfamate did not deamminate; decreases in nitrite ion occurred only as the N02 was oxidized back to nitrate. This was quite evident in the blank and in some of the other tests. Oxidation occurred because of traces of oxygen in the house nitrogen used to blanket the reactions, or possibly because air was introduced when the 2k hour samples were taken. The really significant result, however, is the total nitrogen balance that resulted at the end of Test 3- Of 1*9 ppm of nitrogen initially present, only 39-2 ppm could be accounted for after US hours. These data lead to the conclusion that Fe++ can reduce nitrate ion directly to either Ng or N20 which can escape from the reactor. Since the actual constitution of the* escaping material has yet to be determined, it is generally referred to in this report as "missing N". The precipitated iron insolubles in these reactions were black, which indicated that the iron end product was not ferric hydroxide, but the mixed Fe -2Fe salt; that is, Fe30^, or the corresponding mixed hydroxide, although the existence of the latter does not seem to have to be definitely established. Thus, a balanced equation for the reduction of nitrate to nitrite should be written as in Equation (3). HgO + 3Fe*+ + NO" •* Fe"H'.2Fe*++ + KO" + 2(OH~) ('3) The denitrification reactions may be as those shown by Equations (U) and (5). ^e^ + 2KO- + jHgO t MFe^e^) + NgO + lO(OlT) (k) 15F6++ + 2NO- + oHgO - 5(Fe++.2Fe+++) + N,, + 12(OlT) (5) 18 ------- TABLE j FERROUS ION REDUCTIOIi OF NO. Test No. 1 2 3 k 5 Conditions Varied Sulfamate added after 2k hrs. No NO". Sulfamate added after 2k hrs. NO" added. Sulfamate added initially. No NO". Sulfamate added initially. NOg added. Blank (No Fe++) . NO' added. Tine, hr. 0 2k 0 2k 0 2k kQ 0 2k kQ 0 2k kQ Analytical Results, ppm NO"-N Uii.3 30.5 33.9 &' 1*9.0 30.2 33.0 20J5 3U.5 50.0 50.9 5^.7 NO--N — 5.0 6~7 k.k 9-7 U-3 k.2 3-9 0.8 3 i -- i.e -- •~ Fixed Conditions: Temperature: t>5° F Atmosphere : Nitrogen pll: 11 NO^-N Concentration: 50 ppm Fe++ Concentration: 3X < ,_ _x ; (Denotes 3 moles of agent per mole of NO,) Sulfamate Concentration: 3X ) •* Catalyst: Cu , 5 ppm 19 ------- Finally, some of the nitrate is reduced to MH3, as in Equation (6). 12Fe+* + NO' + ohVjO -> UCFe^-ZPe***) + NH3 + 9 (OH~) (6) These latter reactions (Eq. Jf, 5, 6) have a much greater iron demand than the simple 2 electron N03/N02 reduction. In fact, if it is assumed that the pissing N in Test 3 evolved as N2, then the reduced products (NOg, N2, and NH3) account for about 8o> of the ferrous ion that was originally present That is, reduction of N03 by Fe~ is a fairly efficient process, b£in this first_screening series there simply was not enough of Fe++ to reduce an the N03 via the U, 5, and 8 electron reactions shown. 4^ J^rification was confirmed in additional screening tests, carried out with the Fe ion concentration increased to eight times that of the NO? ion on a molar basis (designated as "8X" concentration). This would be slightly more than enough ferrous ion to convert all the N0§ to N2, assuming N2 to be the denitrification product. In this series, pH and catalyst level were varied. No deammination agent was present in this or any subsequent studies. The data in Table k clearly show that as much as 50-55$ denitrification can occur as a result of direct Fe++ reduction of N03. The data also show that denitrification does not occur under initially acid conditions The effects of catalyst level and pH are not clearly separated, however, and no ready explanation can be given for the reversal of the NO^-N and NOo-N levels shown in the "duplicate" 3a and 3b tests. Attempted Identification of "Missing N" +v. ,,P?11?wing,,tJle above screening tests, an attempt was made to identify the "missing N" by carrying out the reaction in a closed, evacuated vessel using carefully degassed solutions, and examining the evolved gases mass ' spectrometrically. This experiment failed because no denitrification occurred, for reasons unknown. The conditions (except for the atmosphere) were the same as those shown for tests 3a and 3b in Table U. However, the mass spectrometer detected only trace amounts of nitrogen and argon (residual air) after trapping out the water vapor, and analyses of the solution showed N03-N, 1.1 ppm; N02-N, 21.3 ppm; and NH3-N, 25.8 ppm. (U8.2 ppm total for a nominal 50 ppm N03-N initial concentration.) Identification of "missing N" was planned for some later stage once the reaction was more fully understood. However, other aspects of the reaction took on a greater priority, and these plans were never carried out. Nonetheless, these results are of importance: they show that occa- sionally, for reasons as yet unknown, the denitrification reaction fails completely. Accordingly, results discussed subsequently in this report must be interpreted with caution since the occasional failure from unknown causes can be a confounding factor. 20 ------- TABLE DENITRIFICATION WITH FERROUS ION (Duplicate Tests) Test No. 1. a b 2. a b 3. a b pH k.O i*.o 7.1 7.1 11.0 11.0 Cu++ ppm 5 5 10 10 5 5 Time, hr. 0 1*8 0 ua 0 2U W 0 2U ua 0 21* 48 0 2U hQ Nitrogen Balance, ppm NO"-N 50 50 50 50 50 25.8 6.1 5U 35.2 6.1 50 3.0 2.6 50 15.3 13-3 NOg-N -- — 0.5 0 16.0 lit.fc 4~3 ^.3 NH.-N -- __ 16~5 19 19-5 19.1 "Missing N" — — 26.2 2U.9 n.5 11.3 Fixed Conditions: Temperature: 35° F Atmosphere: Nitrogen NO~-N Concentration: 50 ppm Fe++ Concentration: 8X 21 ------- As a final check, the possibility of "missing N" being removed from the system along with the precipitated iron insolubles should be tested, even though this mode of denitrification seems very unlikely. There are no known insoluble nitrates or nitrites that can form in this system, and while metal oxides and their gels (e.g., alumina, Ref. A-50, A-6o) have been shown to absorb NOg and NOjj ions, such absorption is extremely inefficient, even in concentrates. The catalyst ions are also undoubtedly incorporated, at least in part, into the insolubles. However, loss of nitrogen as a cupric ammine, such as Cu(NH3)ii , is also unlikely. For instance, in tests number 2a and 2b (Table U) the molar ratio of "missing N" to cupric ion approaches 12 to 1. Thus, much more denitrification occurs than can be accounted for by the formation of a cupric ammine. Also, as win be discussed subsequently, the pH drops as the reaction progresses. With an initial pH near 7 (as in tests 2a and 2b), the final pH will be near 5. The simple ammine complexes dissociate readily under acid conditions. There are a few known complexes of N20 with salts, such as KgSOVNgO (Ref. A-U). These behave much like the ammines; the NgO is liberated from such species in dilute acids. Thus, although complexes of ammonia or N20 might well play a role in the mechanism of the reaction during its initial alkaline stage, it does not appear that any significant part of the denitri- fication can be attributed to complex formation and loss of such complexes as part of the insoluble materials. Effect of pH on Denitrification An results reported in this section and the remainder of this report were obtained at a 10 ppm NO§-N concentration, in order to generate data at a level near that of the average secondary effluent. The effect of varying pH is shown in Table 5. There is no readily evident correlation of pH and denitrification from these data, even excluding the apparently anomalous results in Test 5a. There is a discernible trend for increased reduction with increased pH, particularly if one considers the extremes (U.3 ppm average residual NO§-N at pH 7.0, 0.2 average residual at pH 11.0). With the exception of Test 7, the predominant reduced products were "missing N" and ammonia. In tests 3 and 6, these two products accounted for 70 to 90$ of the total nitrogen. A combination of processes, where ferrous ion denitrification would be fonowed by ammonia stripping, thus has the potential for close to complete denitrification. It should be noted that pH was not constant during the course of the tests reported in Table 5. The reaction system became more acid as the reduction progressed. Starting at pH 7, the final pH was between 5 and 6. 22 ------- TABLE EFFECT OF pH ON DEIIITRIFICATIOW Test No. 1. a b 2. a b 3. a b k. a b 5. a b 6. a b 7. a b Initial PH 6.0 6.0 7.0 7.0 7.5 7.5 8.G 0.0 8.5 8.5 9.0 9-0 11.0 11.0 Nitrogen Balance, pp:.i NO"-N 9.5 9-5 3.6 2.1 2.0 3-5 3.3 7.0 3.5 0.0 2.3 0.1 0.2 NO^-N 0 0 0.6 0.7 0 0.9 l.ci . 0.9 2.7 1.1 0.1 O.U 6.1 5.7 NH.,-N 0.5 3.0 2.2 3^ 3.5 3.6 0.5 h.6 k.Q 1.1 1.0 Fixed Conditions: Initial NOl-N Concentration: 10 ppm Fe Concentration: CX Catalyst: Cu , 5 ppm Reaction Time: 2U hr Temperature : 85 F Atmosphere: Nitrogen Denitrification, % 0 0 28 22 1*5 37 12 22 0 20 33 27 31 23 ------- Starting at 11, the final pH was between 8 and 9. An additional experiment was carried out in which the pH was readjusted manually during the course of duplicate runs in order to hold it near a value of 7. With all other fixed conditions the same as those shown in Table 5, the final nitrogen balances were as follows: N03-N, 6.9, 5-7 ppm; N02-H, 1.8, 2.3 ppm; WH3-N, 0.8, l.U ppm. Thus, % or less denitrification occurred in these runs where pH was held relatively constant. Effect of Catalysts on Denitrification The results of a series of tests where catalysts and their concentrations were varied are given in Table 6. Silver ion (Ag ) appears to be a much more effective catalyst for reduc- tion than is cupric ion (Cu++) but essentially no denitrification occurred when silver ion was the catalyst. Silver and cupric ions are the classical homogeneous reduction catalysts (See p. Ik for discussion and references) and their effect is reportedly enhanced by the presence of organic acid anions. In subsequent sections of this report, the use of silver and cupric acetate will be noted, but no effect on denitrification can be attributed to this variation. Zinc, cadmium, and magnesium ions had no catalytic effect in the reduc- tion reaction in the absence of organic acids. The ability of Zn++, Cd"1"1", and Ms"1"*" to catalyze the reduction or denitrification in the presence of an organic acid was not examined because of time limitations. It is important to remember that this reaction occurs in a heterogen- eous system, with precipitated ferrous hydroxide being the reducing agent. The catalysts, Cu+* and Ag+ are, however, the classical homogeneous catalysts, and their mode of action under the present conditions is not readily evident. The heterogeneity of this system may be the source of much of the seemingly erratic behavior in this reaction that is now unexplained. Minor variations in the way Pe(OH>2 is precipitated, the extent to which catalyst ions are absorbed in the precipitate, the manner in which the black ferrous-ferric complex forms from the Fe(OH)2, and a number of other factors associated with the solid state could all play a role. Effect of Air on Reduction and Denitrification As noted earlier, the bubbling of air through solutions had confounded our first test results, and data were also presented showing reoxidation of nitrite to nitrate in screening tests (page 19). Analytical studies on this latter effect, which are reported subsequently, showed that nitrite reoxida- tion would not occur if the solutions remained basic. To further test the effect of air, a reaction was conducted at high pH in an open vessel, with the reaction mixtured stirred mechanically. Under these conditions, 100& 2k ------- TABLE 6 EFFECT OP CATALYSTS ON DENITRIFICATION Catalyst Type Cu++ Cu++ Cu++ AE+ A6+ AC+ Fixed Cone . , ppm 1 5 10 1 5 10 Conditions : Ag Nitrogen Balance, ppm NO~-N 6.7 0.2 0.3 0.1 0.3 0.2 NO--N 2.8 5-9 5-9 5-9 8.5 5-3 NH -N 0.8 1.0 — 3.0 l.l 5.1 Denitrification, % None 29 — (Trace?) None None and Cu as the sulfate and chloride NO"-N Concentration: 10 ppm Fe Concentration: 8X pH: 11 (initial) Reaction Time: 2U hr Temperature : 65 F A tmo sphere : Ni trogen Values reported are averages of duplicate runs 25 ------- reduction was observed, which shows that no N0§ is formed by H0j> reoxidation under such conditions. However, no denitrification occurred in this test: the final nitrogen balance showed no nitrate; N02-N, 7.8 ppm, and HH3-N, 2.3 ppm (all fixed conditions, except atmosphere, were the same as Test 7 in Table 5). It is difficult to attribute a failure to denitrify to the presence of air, particularly when reduction has been quantitative. The result may well be confounded by the "occasional unknown cause" mentioned earlier. Lime for pH Adjustment In all the work described earlier in this report, sodium hydroxide was used to make initial pH adjustments. Table 7 shows the effect of using lime (CaO) in place of the NaOH. Lime obviously does not interfere with denitrification unless, as in Test 3, enough lime is added so that the pH remains constant as the lime slowly dissolves. However, some decrease in total reduction may have occurred. Compare Tests la and Ib (average 68% reduction) in Table 7 with Test 6 in Table 5- In the latter test, with NaOH, an average of 85# reduction occurred. Also shown in Table 7 is the effect of increased Fe^level. In Test 2, this drove the reduction largely to MH3, although some denitrification still occurred. Effect of Phosphate and Carbonate Ions The addition of a mixture of phosphate and carbonate ions (as KgCOs and KHgPOli) inhibits the reduction reaction. As the duplicate results in Table 8 indicate, this inhibition is caused by the phosphate, probably by deactivating the catalyst. Phosphate appears to be exceptionally powerful in its ability to inhibit catalysis by Cu4*. In additional tests, increases in Cu1"*" concentration to 20 ppm or a change from CuCl2 to cupric acetate (CuAcg), again with an increase to 20 ppm, were unable to overcame the effect. However, two other methods were tested which did prove effective: increasing the iron concentration and sub- stituting Ag+ for CVL++ as the catalyst. Results are shown in Table 9. Bbte that in all of the tests above, where 003 alone was present, or where the effect of P01|~3 was overcome by one means or another, there was essentially no denitrification, even though reduction reached the 90$+ level. The carbonate ion buffers these reactions, so that pH remains relatively constant throughout the reaction. Similar results (reduction without denitri- fication) in the absence of carbonate were noted earlier, when NaOH or lime was added periodically during the course of a reaction so that the pH remained at a high level. Thus, there is an apparent correlation showing that denitri- fication is suppressed by having a constant pH, although this relationship is partially confounded by the fact that Ag+, which catalyzed some of the buffered runs, does not appear to induce denitrification even in the absence of carbonate. 26 ------- TABLE 7 DENITRIFICATION IN THE PRESENCE OF LIME Test No. 1. a b 2. 3. Lime Added ppm 300 300 700 7^0 _ -H- Fe Level 8X 8X i6x i6x pH Initial 9.0 9.^ 10.2 10.5 Final 6.0 6.1* C.5 10.7 Nitrogen Balance, ppm NO"-N 3.1 3.3 0 h.o NOg-N 1.6 l.U 0.3 2.U NH3-N 3.0 3.0 7.9 3.6 "Mi s sing ' N" 2.3 2.3 1.0 None Fixed Conditions: NOl Concentration: 10 ppm ++ Catalyst: Cu , 5-10 ppm Tempera ture : 6"5° F Atmosphere: Nitrogen Reaction Time: 2b hr 27 ------- TABLE 8 EFFECT OF PHOSPHATE AND CARBONATE IONS pH Adjusted with NaOH NaOH NaOH NaOH Lime Lime NaOH NaOH NaOH NaOH P Initial 9.5-10 9-5-10 n 11 7.0 7.1 10.0 9.1 10.3 9.0 H Final ~ • m 7.5 7.5 9-9 9.7 10.2 8.1* «V3-p, ppm 10 10 10 10 10 10 10 10 None None C03 ppm 100 100 100 100 100 100 None None 100 100 Fixed Conditions: Nitrogen Balance, ppm NO~-N 8.2 7.2 9.3 8.0 7.8 8.1 9.8 8.7 1.8 3.2 NO'-N 1.6 l.U 1.1 l.U 0 0 0 0 U.3 1.8 NH.-N 0.8 0.8 0.3 1.3 1.3 0.7 1.1 3.5 k.k N0§ Concentration: 10 ppm Fe** Level: 6X Temperature: 85° F Atmosphere: Nitrogen Reaction Time: 2k hr Catalyst: Cu'H', 5 ppm 28 ------- TABLE V OVERCOMING THE EFFECT OF PHOSPHATE AND CARBONATE Fe Level l6x 2Ux Ox 8x 5x Catalyst Cu , 5 ppm r. •*-«• c LU , ) ppm Ag , 5 ppm Ag+, 5 ppm AS+» 5 ppm Ag+, 5 ppm (AgAc) t>H Adjusted with NaOH NaOH Lime Lime Lime Lime pH Initial O.Q 9-9 10.5 10.5 10.5 10. I* Final 10.2 10.3 10.3 10.3 9.U 9-1* Nitrogen Balance, ppn NO^-N U.7 3.0 1.7 2.0 0.9 1.1 NO'-N 3A 3.3 6. ii 6.8 5-0 5.3 I«3-N 2.1 u.o 1.2 l.U 3.7 3.3 Fixed Conditions: NO~-N Concentration: 10 ppm PO^-P Concentration: 10 ppm COlr Concentration: 100 ppm Temperature : 85 F Atmosphere: Nitrogen Reaction Time: 2k hr 29 ------- There is no immediately evident explanation for the above effect of constant pH. Obviously, pH-related phenomena win have to be investigated in much more detail before applying this reaction in an actual water reclamation process. ANALYTICAL SUPPORT Background and Selection of Methods Nitrate ion, which represents the most highly oxidized phase in the nitrogen cycle, generally occurs in trace quantities in surface water supplies. Since a limit of 1*5 mg/1 nitrate (10 ppm NOo-N) has been imposed on drinking waters, a number of analytical chemistry methods are available in the literature to determine NO§-N at the ppm level. Methods available in the literature for nitrate ion determination were examined and Judged for relative merit with the particular parameters of the nitrate ion reduction effort in mind. For example, for any given set of experiments, analyses were also required for nitrite ion and to complete a material balance, a final analysis for ammonia might be required. All of these were to be deter- mined in the presence of a reducing agent and possibly a deammination agent. Accordingly, methods for nitrite ion and ammonia were also examined. A brief review of the methods considered is given below. Nitrate and nitrite analyses can be performed according to the method of Fisher, Ibert, and Beckman (Hef. F-6) which utilizes the sulfur-yellow color produced by brucine in sulfuric acid solution. By varying the concentration of sulfuric acid (less than 25$ for nitrites, greater than 5056 for nitrates) the two can be determined on aliquots as small as 15 ml containing 0-1 micro- grams of the anion. A modification of this procedure has been incorporated as a standard method for nitrate in the Standard Methods for the Examination of Water and Wastewater (Ref. F-l). Ferrous and ferric ion have been reported to give slight positive interferences. Interference due to nitrite ion is eliminated by the use of sulfanilic acid. Where nitrate ion alone is desired another common practice is to destroy the nitrite ion using solid sulfamic acid (NH2S03H). The reaction with sulfamic acid (Ref. F-5) is almost instan- taneous and is not interfered with by the species used in the nitrate reduc- tion studies. Nitrite ion may also be determined by the coupling of diazotized sulfanilic acid with 1-naphthylamine hydrochloride at pH 2.0-2.5 with the formation of the reddish purple azo dye (Ref. F-l). The method is sensitive to 0.1 ppm NOg-nitrogen in a 10 ml sample. Bastion, et al., (Ref. F-3) reported an ultraviolet spectrophotometric method for the determination of nitrate ion in alkaline earth carbonates. The method is based on the absorption of nitrate ion in the 200-220 millimicron region. The absorp- tion maximum is 200 nM, but in the systems studied, measurements at 210were found to give optimum results. Armstrong (Ref. F-2) used a modification of this method to determine nitrate in sea water. The samples are run in 50$ 30 ------- and 0.05 M HC1. At these concentrations, both nitrite and nitrate have absorption naxima at 22? mu. An ultraviolet spectrophotometric method has been recommended in the Standard Methods of Water Examination as useful for screening large numbers of drinking water samples for nitrate ion. The speed, accuracy, precision, and interferences of each of the above methods were considered in the selection of a candidate procedure for nitrite and nitrate. Of the methods reviewed, the direct measurement of ultraviolet absorbance appeared most desirable with respect to speed and simplicity and was selected for laboratory evaluation. Methods for determination of ammonia were also examined (Ref. F-l, pp 186-19U). The method of choice was distil- lation from a strong base in a micro-Kjeldahl distillation flask, followed by titration with a standard mineral acid. Since urea was a deammination agent under consideration and consequently a source of nitrogen, methods for its determination at the ppm level were reviewed. Two possible methods appeared appropriate (Ref. F-7, F-9). The former is a spectrophotometric procedure using p-dimethylaminobenzaldehyde and the latter is a conversion to ammonia by enzymatic hydrolysis. Analytical Program The major part of the analytical program consisted of the evaluation and development of the ultraviolet spectrophotometric method for nitrate and nitrite ion by thoroughly checking out possible interferences from the candi- date reducing agents and deammination agents. The ammonia distillation method was examined to determine sensitivity limits and interferences. Subsequent to this, the tentative procedures were applied to the experimental program. From time to tine, analytical anomalies arose as a result of changes in the experimental plan (new reducing agent-deammination agent pair, new catalysts, different pH, etc.), or interferences that were not considered at the begin- ning of the program. These anomalies were investigated and modification or improvements in the methods were made. As the program progressed and ferrous ion became the most frequently used reducing agent, the analysis scheme became routine and the analytical efforts consisted entirely of analysis of samples. A complete description of the ultraviolet spectrophotometric method used is shown in the Appendix at the end of this report. Development of the Ultraviolet Method Some of the reagents selected to be screened for the reduction of nitrate and subsequent deammination of nitrite were procured, and solutions of each were prepared. Solutions of potassium nitrate and potassium nitrite were also prepared. Each solution was subjected to ultraviolet spectrophotometric scanning with a Gary I1* Recording Spectrophotometer, covering the region in which nitrate is absorbant. The results of this examination are shown in Table 10. 31 ------- TABLE 10 UV ABSOKBANCE OP NITRATE, NITRITE, AND CANDIDATE AGENTS Compound Potassium Nitrate Potassium Nitrite Sucrose Urea Formaldehyde fydrazine Sulfate Sulfamic Acid Nitrite + Sulfamic Acid Concentration ppm 1 (as N) 2 (as N) 100 100 100 100 (as Ngfy) 2 (as N) + large excess Wavelength of Abs. Maximum, mia 205 210 < 190 < 190 < 190 < 190 < 190 < 190 Specific Absorbance, = Optical Density g/1 x Cell Thickness 690 375 Not applicable Not applicable Not applicable Not applicable Not applicable Only sulfamic acid absorbance seen 32 ------- Examination of the data in Table 10 reveals that nitrate has an absorp- tion maximum at 205 EU» and nitrite has a maximum at 210 nti, in good agreement with ultraviolet spectra recorded elsewhere (Ref. F-2, F-3). It should be noted that although Standard Methods (Rcf. F-l) cans for measurements at 220 nil, this is not an absorption maximum. The absorbances of the candidate reducing agents at wavelengths below, but near, 190 flJ apparently do not interfere with the nitrate or nitrite absorbance measurements. The solutions vere so concentrated relative to the nitrate and nitrite solutions studied that there was some overlap of peak shoulders (or tail) with the 200 rau region. It is estimated from these studies of individual solutions that at reducing agent concentrations of 50 times that of the nitrate ion, the overlap would introduce an error of about yjt in the nitrate ion determination. The addition of solid sulfamic acid to a 2 ppm nitrite-N solution, fol- lowed by immediate spectrophotometrie scan, was found to destroy quantita- tively and immediately the nitrite ion. This experiment indicated nitrite ion can be removed easily and quickly. Absorbance measurements before and after nitrite removal will yield values for both ions, by difference calcu- lation. For the initial phase of the program, four reducing agents vere chosen for study: J.A Ferrous ion (Fe ) as FeSOU Sulfur dioxide (802) Filtrasorb kOQ carbon Formaldehyde The UV spectrophotometric method was further examined for interferences of the above species at stoichiometric and threefold excess. None of these species seemed to interfere with the nitrate determination. Sulfamic acid at threefold concentrations introduced a small error, but this proved easy to correct by use of a reagent blank. Calibration curves were prepared from stock sodium nitrate and sodium nitrite solutions, respectively, and the method was used to analyze the first experimental mixtures. The analytical procedure involved the measurement of the UV absorbance of a diluted portion of the sample, addition of solid sulfamic acid to the diluted solution, and a re-examination of the UV absorbance. Experience showed that measurements at 200 mu gave the most reproducible results. Since sulfamic acid reacts quantitatively with nitrite ion, the UV absorbance after addition of sulfamic acid was a measure of the nitrate ion and the difference in UV absorbance before and after sulfamic addition was a measure of the nitrite ion. During the course of the early analyses, deammination evidently was not accomplished in some cases in which it was expected. Study of the problem revealed that incomplete deammination appeared to be restricted to cases 33 ------- in which the pH was not highly acidic. Data from experiments performed on known sodium nitrite solutions containing sulfamate ion at various concentra- tions are given in Table 11. The results shown in Table 11 indicated that deammination did not occur rapidly until the solution was quite acid, and that neither nitrate nor nitrite can be measured by UV absorbance in strongly basic solutions because the hydroxyl ion interferes. Basic solutions were found not to deanuninate com- pletely in 1*8 hours, but in acid solution the deammination was very rapid. Indication of the interference by hydroxyl ion prompted pH-absorbance studies. Solutions adjusted to cover the pH range were prepared and their UV spectra were taken on the Gary, Model lU, spectrophotometer. The absor- bances at 210 uu of the solutions are shown in Table 12. It appears that the hydroxyl ion precludes ultraviolet analysis of basic solutions, and the analytical procedure was altered to include addition of excess acid to the aliquot of sample being diluted for analysis. The work of Bastion (Ref. F-3) showed that perchlorate ion shows no UV absorbance at 230-200 muj therefore, acidification of samples was accomplished with perch- loric acid. Concentrated perchloric acid is a powerful oxidizing agent, but when diluted to less than 20$, HC10U has virtually no oxidizing power (Ref. F-8) The amount of perchloric acid necessary to bring the pH of sample solutions below 3 is insufficient to reoxidize nitrite to nitrate during the course of the analysis. Investigation of sulfite ion interference with UV nitrate analysis was undertaken when test results of solutions containing sodium sulfite showed anomalies between sample and reagent blank when sulfamic acid was added to such solutions. The difficulty was resolved when, in the light of the pH studies described above, it was realized that the addition of sulfamic acid to the sample solution during analysis dramatically changed the pH, and the sample and reagent blank may have been buffered more or less by the sulfite ion. The possibility of sulfamic acid decomposition in basic solution to yield ammonium ion was also considered, but since the high pH solutions could not be analyzed as received, this possibility was not pursued. The studies indicated that caution must be exercised in interpreting analytical results of some reduction experiments, since deammination was found to be pH-dependent. The effect of pH changes during the analytical procedure which increase the deammination rate were borne in mind, to prevent alteration of nitrite ion content during the analysis. An examination was made of the air oxidation of standard solutions con- taining varying concentrations of nitrate and nitrite ion. Using standard handling and transfer procedures, reoxidation of significant amounts of N0§ occurred at pH's of 3 or lower when the time of handling exceeded one hour. When analyzing a series of six or more samples, the total analysis time frequently exceeded one hour. At pH's of 7 or above, oxidation did not ------- TABLE 11 ANALYTICAL STUDY OF SULFAMATE DEAMMINATION (3 ppm NO~-N Level) Solution No. 1. a b c d 2. 3. k. a b Treatment No sulfamic acid Increment of sulfamic acid added, lowering pH Further increment added Further increment added Near stoichiometric amount of sulfamic acid added Sulfaraate added, base added Sulfamate, base added Solution acidified pH H.O 3.5 3.0 3.0 2.3 11 6.9 2.7 Absorbance at 210 nu 1.1U 0.82 0.25 0.02 0.05 OH" ion interference 1.12 0.10 35 ------- TABLE 12 EFFECT OF pH ON UV ABSORPTION Solution NaOH HC1 HC10U NaOH + HC1 PH 11 5 3 3.8 Absorbance at 210 ny very strong nil nil 0.03 36 ------- occur during this time period. Therefore, during analyses of ferrous ion reduction tests, oxidation did not occur during the filtration step as long as the solution was above pH 7. To prevent any reoxidation of N0j> at low pH's, the procedure was modified so that each sample was scanned in the appropriate UV range immediately after the acidification step. The procedure described in the Appendix was used to carry out a series of tests on standard nitrate-nitrite mixtures to establish error limits. For samples containing 50 ppm total K, the limits of the procedure were found to be less than 1 ppm for N0§, and less than 2 ppm for NOJ3 when a large amount of NO" is also present. Ammonia Distillation Method Analysis for ammonia was examined to ascertain the lower limits of detection by the distillation method, and possible interferences from nitro- geneous reducing and deammination agents. Some literature information on the stability of the sulfamate ion was obtained (Ref. F-lt); however, analysis of known standards containing ferrous ion and sulfamic acid was considered necessary. The distillation method, deemed most reliable for the purpose of the reduction study, involved the addition of an aliquot of the sample and a volume of sodium hydroxide to a small still. Ammonia in the sample is steam distilled, caught in a boric acid solution, and titrated with dilute standard acid solution. The examination of nitrate and reducing agent solu- tions was performed with 25 ml aliquots of each solution studied. The results are shown in Table 13. The data of Table 13 show that recovery of ammonia from dilute solutions of ammonium ion is sufficiently quantitative to monitor formation of ammonia from nitrate reduction at the levels of interest in this effort. However, the recovery of ammonia from solutions prepared with ferrous sulfate and nitrate, and with sulfamic acid, are potential sources of error. The sul- famic acid contribution of ammonia may be the result of slow decomposition of the reagent. The ferrous sulfate or nitrate alone yielded no ammonia, but the combination yielded ammonia, indicating that reduction of nitrate to ammonia occurred to some extent under the conditions of the distillation. The error, however, was only about 1 ppm at the 50 ppm nitrate nitrogen level and the method was used without further modification. 37 ------- TABLE 13 AMMONIA DETERMINATIONS ON SYNTHETIC NITRATE REDUCTION SAMPLES (25 ml Aliquots in All Cases) Solution Ammonjun salt AmmonJiB salt 50 ppm NOo-N (plus 3 x stoichiometric amount of FeSOVTIfeO (150 mg) plus 3 x stoichiometric amount of NILSO-H (26 mg) 50 ppm NO~-N 150 mg FeSO^'THgO 26 mg NHgSO^H 50 ppm NOl-N plus 150 mg FeSO. 'TILO 50 ppm NO"-N plus 200 mg FeSOr -THpO Ammonia Nitrogen ppm Added 1 10 0 0 0 0 0 0 ppm Recovered 0.96 9-9 1.2 0 0 0.2 0.9 1.9 38 ------- EXPERIMEMTAL DETAILS APPARATUS Reactions were conducted in 200 ml volumetric flasks, each containing a mechanical stirrer and a tube inserted into the neck to provide the appropriate gaseous atmosphere. (During the first few reactions, the gas tubes were in- serted into the liquid, and the gas bubbled through, but this process did not effectively stir suspended solids, and was abandoned.) The reactors were immersed in a thernostatted water bath during the reaction period. Adjustments and changes in pH were monitored with a Beckman pH meter (Model G). REAGENTS Reagents and their sources are listed below: Water, lUO Potassium Nitrate, KNO., Sodium Nitrite, NaN02 Sodium Hydroxide, NaOH Ferrous Sulfate, FeSOj/THpO Carbon, C Sulfur Dioxide, S02 Formaldehyde, CHgO Glucose, CgHipOg Iron Powder, Fe Hydrazine, NpH. Hydrazine Sulfate, NpHgSCV Carbon Monoxide, CO Sulfamic Acid, NHgSO_H Cupric Chloride, CuCl2 Silver Sulfate, AggSO^ Cupric Acetate, Cu(CpH Ogjp Silver Acetate, AgCgH,02 - From a laboratory demineralizer - J. T. Baker Chem. Co. - J. T. Baker Chem. Co. - Mallinckrodt Chem. Works - J. T. Baker Chem. Co. - Calgon Corporation's Filtrasorb - The Matheson Co. - 36.6% Formalin, J. T. Baker Chem. Co. - Eastman Organic Chem. Co. - Mallinckrodt - Eastman - Fisher Scientific Co. - The Matheson Co. - Eastman - Mallinckrodt - J. T. Baker - J. T. Baker - Mallinckrodt 39 ------- Ferric Chloride, Ammonium Vanadate, Utt^MO Vanadium Pentoxide, V^ Lime» CaO Potassium Carbonate, KgCO- Potassium Phosphate, Monobasic, - J. T. Baker - Harshaw Scientific - J. T. Baker . j. T. Baker - J. T. Baker . General Chem. Division, Allied Chem. PROCEDURE Note ha an Note that a blank the redUCing *** is Ascribed below. water with no NO^) was treated in the same J *««•«» m me same manner as described for the sample. »/"?? ^les were PreP«ed by adding standard nitrate solutions (2 mg }y meanS °f a Sma11 burette to water contained in 200 ml volumetric AhS™* f^hC Wa?6r *' initiaU^ Just short <* the desired 200 ml. After the addition of the various reagents, the volume was adjusted to the 200 i??J'«re£OUV0n WaS addCd t0 the nitrate solution in ^ form of the s'Siton ?^™2? (c^S??^Wei8tad °n the •M^1C1 f*1^). The catalyst solution (a*., 1 ng Cu++/nil) was then added by means of a pipette. After thorough mixing, the pH of the solution was adjusted to the desired level by adding a few drops of a sodium hydroxide solution (6M). The pH of the blank was first adjusted in order to establish the volume of base necessary to reach the desired pH. Addition of the same volume of base to the test samples, made further adjustments unnecessary. Next, the flasks were Blaced in a constant temperature bath (85° F) and siirred m^chanicalS unler a blanket of nitrogen gas. At the end of the reaction period, samples^ wUhL-awn for pH detennination, nitrate-nitrite analyses and ammonia analysis. "&Wn Ior Deammination experiments involved urea and sulfamic acid. These were usually added Just prior to the pH adjustment. Urea was added as the solid ' For the experiments where phosphate, carbonate, and lime were used the order of addition consisted of: adding phosphate solution (2 mg P/ml) to the nitrate solution, then adding the carbonate solution (20 mg col/ml) adding weighed amounts of lime (with thorough shaking), then add^g FlsSj and ^ adding the catalyst solution at the very end ------- TEST RESULTS With the exception of the reducing agent screening series, details of all reactions have been tabulated in the DISCUSSION AND RESULTS section of this report. The screening test results are given here in Table lU. Ul ------- TABLE SCREENING OF REDUCING AGENTS Fixed Conditions: NO" Concentration: Temperature: Atmosphere: Reaction Time: Reducing Agent Concentration: 50 ppm 85° F Nitrogen 1*8 hr 3x J Denotes 3 moles of agent per Deaomination Agent Concentration: 3x ) mole of NO" (Occasional variation in conditions noted in test) Reducing Agent so2 so2 so2 so2 Carbon C C C C Deammi nation Agent (when added) NaSO-NHg NaSOoNHo (21* fir) HS03NH2 (start) HSOoNHp (start) HS03NHp (2U hr) HS03NH2 (2U hr) HSOoNHg (Start) HS03NH2 (Start) Catalyst (concentration, ppm] Fe~+(2.5) Cu++ (2.5) Fe+"(2.5) Cu++(2.5) Fe+++(5) Cu^(5) Fe+"<5) Cu++(5) PH 9 9 6 6 6 6 6 6 NO^-N, ppm After 2k hours ^9-7 51.8 50.7 50.2 50.8 51.1* 1*9.2 50.9 After 1*8 hours 1*8.1* 1*9.9 50.8 50.7 51.1 52.1* 50.3 53.1* (?) 1*2 ------- TABLE 1^4 (CONT'D) Reduci:ig Agent Formaldehyde CH 0 2 CH.O CHgO CHgO Carbon Monoxide CO CO Glucose C6H12°6 C6H12°6 C6H12°6 C6H12°6 C6H12°6 Dearami nation Agent (when added) HSOoNHj (stSrtT- HSO^NH? ( start ) HS03NH2 (2k hr) HSOoNHg (21; hr) None «<«,>, CO(NH2)2 (2k hr) CO(NH2)2 (2U hr) CO(NH2)2 (start) None None Catalyst (concentration, ppm) Fe+++(2.5) Cu++(2.5) Fe+~(5) Cu++(5) Cu++(5) Cu++(5) Fe+++(5) Cu++(5) V+5(5) V+5(10) v+5(io) pH 6 6 6 6 6 6 6 6 6 1* k ]1 NO"-N, ppm After 2k hours U9.6 50.2 50.2 50.9 U9-8 51.7 ^9-5 U9.U 50.1 _ kj.6 (NOg-N,*3.5) After U8 hours 1*9.9 ^9-7 50.7 51.2 -- 1*9.0 i.9.1 U8.5 (NOg-N, 2.8) U8.U (NOg-N, l.U) U9.8 (NO'-N, 2.6) ------- TABLE Ik (CONT'D) Reducing Agent Hydrazine Ngfy Ngfy NgH. Hydrazine Sulfat S» NoH/rSOi, (3.75XJ ^HgSOij. (3.75X) NoIkSO), (3.75X) Iron Powder Fe Fe Fe Fe (U.5X) Fe Deammination A .A. Agent (when added) CO(NH2)2 (21* hr) CO(NH2)2 (21* hr) CO(NH2)2 (start) e None None None None HSOdnu (21* hr} HS03NHo CO(NHg)2 (start) None None Catalyst (concentration, ppm) Fe+"H"(0.7) Cu**(0.7) Cu++(0.5) Cu+*(0.25) Cu++(0.25) Cu (l) Cu++(l) Fe++(2.5) Cu^(2.5) Cu++(2.5) Cu+*(lo) » » Cu++(lO) PH n n n 11 n n 11 6 6 6 6 6 NO"-N, ppm After 2k hours 50.1* J+6.2 • 1*8.8 _ 1*8.2 1*3.9 (NOg-N, 8.6) 1*8.6 (NOg-N, 3-2) 1*8.0 (NO'-N, !*.!*) 53.6(7) 1*5.3 !*9.5 1*0.1* (NO'-N, 6.8) 29-9 (NOg-N, ll*.8 After 1*8 hours M.5 1*3.6 30.1* (NO'-N, 0.1*) 1*0.1* (NOg-N, 0) 1*1.8 (NOg-N, 5.2) Ul.3 (NO--N, 5.U) 59.5(?) 1*3.5 1*8.9 • 27.3 (NOg-N, U.8) 27.1* (NO'-N, 10.3) kk ------- REFERENCES A. NITRATES IN WATER General Chemistry 1. Jolly, W. L., The Inorganic Chemistry of Nitrogen. W. A. Benjamin Co., New York, New York, 2. Mbeller, T., Inorganic Chemistry. John Wiley & Sons, New York, New York. 1952. 3. Szabo, Z. G., and Bartha, L. G., Recent Aspects of the Inorganic Chemistry of Nitrogen. 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Key, A., and Etheridge, W., "Further Studies in the Treatment of Gas- Works Liquors in Admixture with Sewage," Inst. Sewage Purif. (Engl.) 1936, Pt. II, 278-300. ------- A. NITRATES IN WATER Natural and Biological Denitrification (Continued) 39- Lockett, W. T., "The Phenomena of Rising Sludge in Relation to the Activated Sludge Process," Surveyor IQk, 37-8 (19^5). kO. McLachlan, J. A., "The Settlement and Rising of Activated Sludge," Surveyor go, 39-UO (1936). *H. Mountfort, L. P., "Some Experiences with Rising Sludge in Humus Tanks." Surveyor iw, 65-6 (19^5). te. O'Shaughnessy, F. R., and Hewitt, C. H., "Phenomena Associated with the Role of Nitrogen in Biological Oxidation," J. Soc. Chem. Ind. 5U. 167-97 (1935). — U3. Parkhurst, J. D., Dryden, F. D., McDermott, G. N., and English, J "Pomona Activated Carbon Pilot Plant," J. Water Pollution Control Federation 32 (10), R70-81 (1967). kh. 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Elimination Processes. Excluding Chemical Reduction (See also Ref. A-36) ^9- Bringmann, G., "Optimal Nitrogen Removal by Addition of Nitrated Living Sludge and Oxidation-Reduction Control," Gesundh.-Ingr. 81, lUO-2 (1960). 50. Calvet, E., Boivinet, P., Noel, M., Thibon, H., Maillard, A., and Tertian, R., "Alumina Gels," Bull. Soc. Chem. France 1953. 99-108. U8 ------- A. NITRATES IN WATER Elimination Processes (Continued) 51. Christenson, C. W., Rex, F. H., Webster, W. M., and Vigil, F. A., "Reduction of Nitrate-N by Modified Activated Sludge," U.S. At. Energy Comm. TID-7517 (pt. la) 26U-7 (1956). 52. Downing, A. L., Tomlinson, T. G., and Truesdale, G. A., "Effect of Inhibitors on Nitrification in the Activated Sludge Process," Inst. Sewage Purif. J. Proc. 196!* (6), 537-51*. 53. Dukes, E. K., and Siddall-III, T. H., "Tetrabutylurea as an Extractant for Nitric Acid and some Actinide Nitrates," J. Inorg. Nucl. Chem. 28 (10), 2307-12 (1966). ~ 5U. Faber, F. M., Olson, H. G., and Taylor, W. A., "What is the Life of Silica Gel?" Chem. Met. Eng. 28, 805 (1923). 55. Farrell, J. B., Stern, G., and Dean, R. B., "Nitrogen Removal from Wastewaters," Envir. Sci. Technol., in press. 56. Fletcher, J. M., and Hardy, C. J., "Extraction of Metal Nitrates by Bu-PO^-HNO " Nucl. Sci. Eng. 16, 1*21-7 (1963). 57. Fresenius, W., Bibo, F. J., and Schneider, W., "Pilot Plant Results on the Removal of Nitrate Ions from Tap Water by Anion Exchangers." Gas Wasserfach 107 (12), 306-9 (1966). 58. Gad, G., "Use of Activated Carbon for Determination of Nitrate, Nitrite and Ammonia in Water and Effluents," Gas-u. Wasserfach 79, 166-7 (1936). 59. Knoch, W., "Extraction of Nitric Acid with Amines," J. Inorg. Nucl. Chem. 27 (9), 2075-91 (1965). 60. Kubli, H., "Information on the Separation of Anions by Adsorption on Alumina," Helv. Chim. Acta 30, 1*53-63 (191*?). • 6l. Ludzack, F. J., and Ettinger, M. B., "Controlling Operation to Minimize Activated Sludge Effluent Nitrogen," J. Water Pollution Control Federa- tion 3{t, 920-31 (1962). 62. Myrick, N., Busch, A. W., and Dawkins, G. S., "Activated Carbon Adsorp- tion, a Unit Process in Liquid Industrial Wastes Treatment," Proc. Ontario Ind. Waste Conf. 10, 193-210 (1963). 63. Reznik, A. M., Potapov, G. G., Korovin, S. S., and Aprakin, I. A., "Extraction of Nitric Acid in the Presence of Sulfuric Acid by Tributyl Phosphate," Zh. Neorgan. Khim. 11 (8), lB^k-6 (1966). ------- A. NITRATES IN WATER Elimination Processes (Continued) 6*. Hohlich, G. A., "Chemical Methods for Removal of Nitrogen and Phosphorous from Sewage-Plant Effluents," Robert A. Taft Sanitary Eng. Center, Tech. Kept, wol-3, 130-5 (1961). 65. Sigworth, E. A., "Potentiality of Active Carton in the Treatment of Industrial Wastes," Proc. Ontario Ind. Waste Conf. 10, 177-92 (1963). 66. Tsitovich, I. K., "Concentration of Ions on Chromatographic Grade Alumina for Qualitative Microanalysis," Tr. Kubansk. Sel'skokhoz Inst. 19P** (9)) 20*1-8. 67. Zeitoun, M. A., Davison, R. R., White, F. B., and Hood, D. W.f "Solvent Extraction of Secondary Waste Water Effluents: Heterogeneous Equilibrium 11 J> Water ponution Contro1 Federa- 68. Zawodna, B., "The Effect of Activated Carbon on the Determination of Nitrogen Compounds in Sewage," Przemysl Spozywczy lU, U30-2 (1960). B. WATER RECLAMATION Conventional and Tertiary Treatment (See Also References A-29 thru A-55) Feasible»" Chem' and En«- News 2. Cooper, R. B., "The Treatment of Waste Water Containing Inorganic Materials," Aust. Chern. Process. Eng. lg (10), 36-8, kO-1 (1966). 3. Gulp, R. L., "Wastewater Reclamation by Tertiary Treatment," J. Water Pollution Control Federation 3£, 799-806 (1963) . k. Dietrich, K. R., "The Third Step in the Purification of Waste Waters to Prevent Eutrophy of Lakes," Chemiker. Ztg. 8? (21), 772-5 (1963). 5. Eckenfelder, W. W. Jr., and O'Conner, D. J., Biological Waste Treatment. Permagon Press, New York, New York, 1961. - - ' 6> £Sf8' °* V*' Water greatMnt; A Guide to the Treatment of Water and Effluents Purification. 3rd Ed. London, Technical Press, 1965. - 7. Malhotra, S. K. , "Nutrient Removal from Secondary Effluent by Alum Ho- 6 50 ------- B. WATER RECLAMATION Conventional and Tertiary Treatment (Continued) 8. Marshall, J. R., "Today's Wastes: Tomorrow's Drinking Water?", Chem. Eng. 6_2, No. 16, 107-10 (1962). 9. Powell, 3. T., "What Part Does Chemical Coagulation Play in Today's Water Treatment Practices?" Power 9_8, No. 1, 80-2, 216, 218, 220 (195*0. 10. Rand, M. C., "General Principles of Chemical Coagulation," Sewage and Ind. Wastes 31, 863-71 (1959). 11. Stephan, D. G., "Water Renovation, Current Status of the Technology," Proc. Southern Water Resources Pollution Control Conf. lU, 113-22 (1965). 12. van Vuuren, L.R.J., Stander, G. J., Henzen, M. R., Me i ring, P.G.J., and van Blerk, S.H.V., "Advanced Purification of Sewage Works Effluent Using a Combined System of Lime Softening and Flotation," Water Research 1, U (1967). Iron Salts in Water Treatment 13. Dean, R. B., "Ultimate Disposal of Wastewater Concentrates to the Environ- ment," Envir. Sci. Technol., in press. lU. Dobrynin, F. T., "The Purification of Water with Iron Vitriol," Vodosnabshenie i Sanit. Tekh. 16, No. 5, 50-3 (19^1). 15. Dodonov, Ya. Ya., and Plekhanova, T. G., "FeSOij. as a Coagulating Agent in the Purification of Water," Vodosnabzhenie Sanit. Tekh. IgUO, No. 12, UO-3; Khim. Referat. Zhur. U, No. 7-8, 98 (19^1). 16. Kunzel-Mehner, A., "Ferric Chloride in the Chemical and Mechanical Purification of Water," Chem. Tech. 15, 129-35 17. Mehner, A. K., "Chemical-Mechanical Treatment of Water with Ferric Chloride," Chem. Tech. 15, 129 (19^2). 18. Pirnie, M., "Some Operating Experiences with Iron and Iron Coagulants in Water Treatment," J. New Engl. Water Works Assoc. 5_1, U37-53 (1937). 19. Schworm, W. B., "Iron Salts for Water and Waste Treatment," Public Works 2jt (10), 118-20 (1963). 20. Scouller, W. D., "Effect of Iron on Sewage Purification," Surveyor 82, (1932). 51 ------- B. WATER RECLAMATION Iron Salts in Water Treatment (Continued) 21. Simmons, P. D., "Ferrous Sulfate as a Coagulant," Proc. 12th Ann. Conf Water Purif., in W. Va. Univ., Eng. Expt. Sta. Tech. Bull. No. 11, 21-3 (1938). 22. Streander, P. B., "Sewage Treatment with Ferrous Sulfate and Aeration " Public Works 6J>, No. 3, 29 (1933). ' 23. Vadyukhim, I. I., "Coagulation of Water with Ferrous Sulfate in Com- bination with Chlorine," Vodosnabzhenie i Sanit. Tekh. lU, No. 10 35-^3 (1939). ' 2U. Wolman, A., "The Role of Iron in the Activated Sludge Process " Ens News Rec. £8, 202-1* (1927). ' *' 25. Zhuchkova, A. M., "Coagulation of Water with Ferrous Sulfate " Teplo- Silovoc Khoz. 1^32, No. 7, 51-2; Khim Referat. Zhur. 1939. No. 12, 86-7. Carbon in Water Treatment (See Also Reference A-U3, -U6, -58, -62, -65, -68) 26. Johnson, R. L., Lowes, F. J. Jr., Smith, R. M., and Powers, T. J., Evaluation of the Use of Activated Carbons and Chemical Regenerants in Treatment of Waste Water," AWTR-11, U.S. Department of Health, Education and Welfare, Public Health Service Publication No. 999-WP-13, Jfey, 196^. 27. Joyce, R. S., and Sukenik, V. A., "Feasibility of Granular Activated- Carbon Adsorption for Waste-Water Renovation, 2," AWTR-15, U. S Department of Health, Education and Welfare, Public Health Service Publication No. 999-WP-28, October, 1965. 28. McGlasson, W. G., Thibodeaux, L. T., and Berger, H. F., "Potential Use Carbon for Waste Water Renovation," Tappi to (12), 521-6 29. Reissaus, K., and Rummel, W., "Use of Activated Carbon in Water Purifi- cation, Wasserwlrt.-Wassertech. 16 (12), U13-16 (1966). 30. Shane. M. S., "How to Black Out Algae," Water Works Eng. 116 (7) 552-T (1963). - ' J C. REDUCING AGENTS FOR NITRATE Ferrous Salts (See Also Reference A-U8, C-79) 1. Abel, E., "Kinetics of the Oxidation of Ferrous Ion by Nitric Acid " Monatsh. 68, 387-93 (1936). ' ------- C. REDUCING AGENTS FOR NITRATE Ferrous Salts (Continued) 2. Abel, E., Schnid. H., and Pollak, F., "Kinetics of the Oxidation of Ferrous Ions by Nitrous Acid," Monatsh. 6_Q, 125-1*3 (1936). 3. Baudisch, 0., and Meyer, P., "The Reduction of Nitrites and Nitrates," Biochem. Z. 107. 1-^2 (1920). I*. Benner, J. M., and Shaw, K., "Reduction of Nitrate by Ferrous Hydroxide Under Various Conditions of Alkalinity," Analyst §0, 626-7 (1955). 5. Brown, L. L., and Drury, J. S., "Nitrogen Isotope Effects in the Reduction of Nitrate, Nitrite, and Hydroxylamine to Ammonia. I. In Sodium Hydroxide Solution with Fe(ll)." J. Chem. Phys. U6 (7), 2833-7 (1967). 6. Carsley, S. H., "The Reduction of Alkali Nitrates by Hydrous Ferrous Oxide," J. Phys. Chem. 3Jt, 176-87 (1930). 7. Gottlieb, 0. H., and Magalhaes, M. T., "The Volumetric Determination of Nitrate Ions," Anal. Chem. 50, 995-7 (1953). 8. Karstein, P., and Grabe, C.A.J., "Determination of Nitrate According to Cotte and Kahane," Chem. Weekblad UU, 237-8 (19^8). 9. Kolthoff, I. M., Sandell, E. B., and Moskovitz, B., "Volumetric Deter- mination of Nitrates with Ferrous Sulfate as Reducing Agent," J. Am. Chem. Soc. 55, lU5**-7 (1933). 10. Krejci, F., and Kacetl, L., "Determination of Nitrate by Titration with Ferrous Sulfate," Chem. and Ind. (London) 1957. 598- 11. Laccetti, M. A., Semel, S., and Roth, M., "Colorimetric Determination of Organic Nitrates and Nitramines," Anal. Chem. 31, 10^9-50 (1959). 12. Miyamoto, S., "The Reducing Action of Ferrous Hydroxide," Japan J. Chem. 1, 57-80 (1922). 13. Murakami, T., "Rapid Volumetric Determination of Nitric Acid or Nitrate by Reduction with Ferrous Salt," Japan Analyst U, 630-3 (1955). I1*. Murakami, T., "Photometric Determination of Nitrite by Using FerroUs Sulfate and Phosphoric Acid," Kagyo Kagaku Zasshi 63, 1295-8 (1960). 15. Pappenhagen, J. M., and Looker. J. J., "Suggested Reduction Methods for the Determination of Nitrates/1 J. Am. Water Works Assoc. £1, 1039-^5 (1959)• 53 ------- C. REDUCING AGENTS FOR NITRATE Ferrous Stilts (Continued) 16. Sandonnini, C., and Bezzi, S., "Reduction of Nitrates with Ferrous Hydroxide," Gazz. Chim. Ital. 6j>, 693-700 (1930). 17. Schoer, E., "Kinetics and Mechanism of the Reaction Between the Ferrous Ion and Nitrous and Nitric Acids," Z. Physik. Chem. A176. 20-1*7 (1936). 18. Szabo, Z. G., and Bartha, L., "Volumetric Determination of Very %aH Quantities of Nitrate," Mikrochemie ver. Mikochim. Acta 38, 1*13-18 (1951). — 19- Szabo, Z. G., and Bartha, L., "A New Titrimetric Method for the Deter- mination of Nitrate Ion," Anal. Chim. Acta £, 33-Uj (1951). 20. Szabo, Z. G., and Bartha, L., "Catalysis in Analytical Chemistry. I. Silver Catalysts in the Reduction of Nitrates by Ferrous Hydroxide " Acta Chim. Huag 1, 116-23 (1951). 21. Szabo, Z. G., and Bartha, L., "Alkalimetric Determination of the Nitrate Ion by Means of a Copper-Catalyzed Reduction," Anal. Chim. Acta 6, U16-1Q (1952). 22. Chao, Tyng Tsair, and Kroontje, Wybe, "Inorganic Nitrogen Transforma- tions Through the Oxidation and Reduction of Iron," Soil Sci. Am Proc. 30 (2), 193-6 (1966). 23. Young, G. K., Bungay, H. R., Brown, L. M., and Parsons, W. A., "Chemical Reduction of Nitrate in Water," J. Water Pollution Control Federation 36, 395-8 (19&). Carbon . 21*. Bylo, Z., and Panek, M., "The Influence of Oxidation on the Reaction of Hard Coals with Dilute Solutions of Nitric Acid," Arch. Gornictwa 9 (k) 383-97 (196U). ^ ^ '' 25. Donnet, J. B., and Lahaye, J., "Oxidation of Carbon Black by Nitric Acid. I. Mode of C02 Formation: Kinetic Aspect of the Reaction," Bun. Soc Chim. France 1966 (If), 1282-5. 26. Farrell, J. B., and Haas, P. A., "Oxidation of Nuclear-Grade Graphite by Nitric Acid and Oxygen," Ind. Eng. Chem., Process Des. Develop. 6 (3) 277-81 (1967). ~ ' 27. van Krevelen, D. W., "A Gas Rich in Nitric Oxide by Reduction of Nitric Acid with Carbon," Brit. 66l, 90U, Nov. 28, 1951. ------- C. REDUCING AGENTS FOR NITRATE Carbon (Continued) 28. Larina, N. K., Khalmukhamedova, R. A., and Tadzhiev, A. T., "Products of Oxidation of the Angrensk Brown Coals by Nitric Acid," Khim. Klassi- fikatsiya Iskop. Uglei, Akad. Nauk. SSSR, Inst. Goryuch. Iskop. 1966 98-107. Sulfur Dioxide 29. Ivin, K. J., "Reaction of Nitrates with Liquid Sulfur Dioxide," Nature 180. 90 (1957). 30. Smedslund, T. H., "Continuous Preparation of Nitric Oxide from Nitric Acid and S02," Finska Kemist-Samfundets Medd. 5£, 37-9 (1950). 31. Soibelman, B. J., and Bresler, F., "Detection of Nitrates in Presence of Interfering Anions," Zavodskaya Lab. £, 359-60 (19^0). 32. Veprek-Siska, J., and Uher, L., "Reduction of the Nitric Acid by Means of Sulfur Dioxide," Collection Czech. Chem. Coramun. 31 (11), U363-71* (1966). 33. Zeegers, R.N.G., "Hydroxylamine Compounds," U.S. 2,555,667, June 5, 1951. Formaldehyde 35. Adams, W. H., Fowler, E. B., and Christenson, C. W., "A Method for Treating Radioactive Nitric Acid Wastes Using Paraformaldehyde," Ind. Eng. Chem. 52, 55-6 (1960). 36. Cultrera, R., and Farrari, E., "Research on the Photochemical Reduction of Nitrate," Ann. Chim (Rome) U7, 1321-36 (1957); W, lUlO-25 (1958); 1*2, 176-82 (1959). 37- Evans, T. F., "Pilot Plant Denitration of Purex Wastes with Formalde- hyde," U.S. At. Energy Comm. HW-58587 (1959). 38. Forsman, R. C., and Oberg, G. C., "HCHO Treatment of Purex Radioactive Waste," U.S. At. Energy Comm. HW-79622 (1963). 39. Halliday, H. M., and Reade, T. H., "Action of Nitrous Acid on Formal- dehyde, " J. Chem. Soc. 19^0. lte-3. UO. Healy, T. V., "The Reaction of Nitric Acid with Formaldehyde and with Formic Acid and its Application to the Removal of Nitric Acid from Mixtures," J. Appl. Chem. 8, 553-61 (1958). Ul. Healy, T. V., "Concentration of Aqueous Metal Salt Solutions Containing Nitric Acid," U.S. 2,835,555, May 20, 1958. 55 ------- C. REDUCING AGENTS FOR NITRATE Formaldehyde (Continued) U2. Kourim, V., and Konecny, C., "Decomposition of Nitric Acid by Formal- dehyde," Chem. Listy 51, 1376-7 (1957). J*3. Morris, J. B., "The Reaction of Nitric Acid with Formaldehyde " Eneraie Nucleaire 1, 216-2U (1957). e UU. Nemtsov, M. S., and Trenke, K. M., "ape stigat ions in the Field of Acid Catalysis. I. Kinetics and Mechanism of the Reactions of Formaldehyde in Acid Aqueous Solutions," Zhur. Obshchei Khim. 22, U15-29 (1952). U5. Nemtsov, M. S., and Trenke, K. M., "Investigations in the Field of Acid Catalysts. I. The Kinetics and Mechanism of the Reactions of Formal- dehyde in Acid Aqueous Solutions," J. Gen. Chem. USSR 22, ^85-96 (1952). U6. Shtol'ts, A. K., "Reaction of Nitric Acid with Formaldehyde, Rongalite and Hydrosulfite," Izvest. Vysshikh Ucheb. Zavedenii, Khim. i Khim. !*7. Vanino, L., and Schinner, A., "The Reaction Between Formaldehyde and Nitrous Acid," Z. Anal. Chem. 5J2, 21-6 (1913). Sugars (See Also Reference C-36) U8. Bray, L. A., and Martin, E. C., "Invention Report - Use of Sugar to Neutralize Nitric Acid Waste Liquors," U.S. At. Energy Conm. HW-75565 (1962) . kg. Bray, L. A., and Martin, E. C., "Removal of Nitric Acid and of Nitrite and Nitrate Ions from Radioactive Waste," U.S. 3,158,577, Nov. 2k, 50. Breit Schneider, R., and Kopriva, B., "Oxidation of Sucrose with Nitric Acid," Listy Cukrovar. 82 (9), 215-20 (1966). 51. Coppinger, F. A., "Pilot Plant Denitration of Purex Water with Sugar " AEC Accession No. 352U3, Rept. No. HW-77080, Avail. OTS (1963). 52. Justat, A., Gorzka, Z., and Janio, K., "Oxidation of D-Glucose with Nitric Acid to Oxalic Acid," Chem. Stosowana 7 (3), 1*09-11* (1963). 53. Kopriva, B., Markova, J., and Breit Schneider, R., "Oxidation of Sucrose with Nitric Acid. II. Kinetics of Oxidation," Listy Cukrov 83 (2) 36-9 (1967). ~ ' 5^. Lesquibe, F., "Degradation of Glucose by Oxidation with Aqueous Acid (HN03) " J. Rech. Centre Natl. Rech. Sci. Lab. Bellevue (Paris) Ik (62) 33-71 (1963). — " 56 ------- C. REDUCING AGENTS FOR NITRATE Sugars (Continued) 55. Soltzberg, S., "Tartaric Acid," U.S. 2,360,196, July 10, 191*5. 56. Tang, Teng-Han, and Kao, F. C., "Preparation of Oxalic Acid I," J. Chem. Eng. China 16, 32 (1939). Powdered Iron 57. Babson, J. A., Burch, W. G. Jr., and Woodis, T. C. Jr., "Critical Evaluation of the Reduced Iron Method for Reduction of Nitrate," J. Assoc. Offic. Agr. Chemists U6 (U), 599-603 (1963). 58. Delius, I., "Removal of Nitrates from Drinking Water," Gesundheits-Ing. 00, 181 (1959). 59. Gehrke, C. W., and Johnson, F. J., "Efficiency of Various Iron Powders in Seducing Nitrate," J. Assoc. Offic. Agr. Chemists J4£, U6-9 (1962). 60. Travers, A., and Diebold, R., "The Action of Nitric Acid on Iron and Iron Carbide (Fe_C)," Bull Soc. Chim. (5), £> 690-3 (1938). 6l. Vetter, K. J., "The Active State and the Spontaneous Repassivation of Current-Activated Iron in Nitric Acid," Z. Electrochem. j>6, 106-15 (1952). Hydrazine and its Salts 62. Bursa, S., and Straszko, J., "Eudiometric Determination of Nitrate Ion in Aqueous Nitric Acid," Chem. Anal. 8, 29-Uo (1963). 63. Davies, A. W., and Taylor, K., "Application of the Auto-Analyzer in a River Authority Laboratory," Technicon Symp., 2nd, N.Y., London 1965, 29U-300 (Pub. 1966). ^^ 6k. Dey, B. B., and Sen, H. K., "Action of Hydrazine Sulfate Upon Nitrites and a New Method for Determining Nitrogen in Nitrites," Z. Anorg. Chem. 71, 236-U2 (1911). 65. Dzhardamalieva, K. K., "Catalytic Reduction with flydrazine," Tr. Inst. Khim. Nauk, Akad. Nauk Kaz.USSH 8, 150-6 (1962). 66. Kahn, L., and Brezenski, F. T., "Determination of Nitrate in Estuarine Waters. Comparison of a Hydrazine Reduction and a Brucine Procedure, and Modification of a Brucine Procedure," Environ. Sci. Technol. 1 (6), U88-91 (1967). 57 ------- C. REDUCING AGENTS FOR NITRATE Hydrazine and its Salts (Continued) 67. Kamphake, L. J., Hannah, S. A., and Cohen, J. M., "Automated Analysis for Nitrate by Hydrazine Reduction," Water Res. 1 (3), 205-16 (1967). 68. Koltunov, V. S., Nikol'akii, V. A., and Azureev, Yu. P., "Kinetics of Hydrazine Oxidation in the Presence of HNOs in an Aqueous Solution " Kinetiki i Kataliz. 3 (6), 077-81 (1962). * ^uwon, 69. Mullin, J. B., and Riley, J. P., "The Spectrophotometric Determination of Nitrate in Natural Waters, with Particular Reference to Seawater " Anal. Chem. Acta 12, k6k-Qo (1955). Miscellaneous (Salts. Metals, etc.) 70. Banerjee, P. J., "Vanadous Sulfate as a Reducing Agent. II. Estimation of Chlorates, Nitrates and Persulfates," J. Indian Chem. Soc. 13, 301-U (1936). — 71. Bartow, E., and Rogers, J. S., "Determination of Nitrates by Reduction with Aluminum," Univ. 111. Bull., W. S. Series 7, lU-27 (1910). 72. Fletcher, J. M., and Woodhead, J. L., "The Reaction of Ruthenium (ill) with Nitric Acid," J. Inorg. Nucl. Chem. 27 (?), 1517-19 (1965). 73. Frank, J. A., and Spence, J. T., "The Reduction of Nitrite by Mo (V)." J. Phys. Chem. 68 (8), 2131-5 (196U). 7b. Gasser, J.K.R., "Substitute Reagent for Titanous Sulfate for Reducing Nitrate Nitrogen," Analyst 88, 237-8 (1963). 75. Guymon, E. Park, and Spence, J. T., "The Reduction of Nitrate by Mo (V) " J. Phys. Chem. 70 (6), 196^-9 (1966). 76. Haight, G. P. Jr., Mohilner, P., and Katz, A., "The Mechanism of the Reduction of Nitrate. I. Stoichiometry of Molybdate-Catalyzed Reductions of Nitrate and Nitrite with Sn (II) in Hydrochloric and Sulfuric Acids " Acta Chem. Scand. 16, 221-8 (1962). 77. Haight, G. P. Jr., and Katz, A., "The Mechanism of Reduction of Nitrate. II. The Kinetics and Mechanism of the Molybdate-Catalyzed Reduction of Nitrate by Sn (II) in Acid Solution," Acta Chem. Scand. 16, 659-72-(1962). 78. Kasbekar, G. S., and Nonnand, A. R., "Reaction Between Nitric Acid and Tin in Presence of Catalysts. II," Proc. Indian Acad. Sci. 1QA. 37.^0 (1939). 79- Milligan, L. H., and Gillette, G. R., "The Reduction of Free Nitric Acid by Means of Ferrous, Stannous or Titanous Salts," J. Phys. Chem. 28. - — 58 ------- C. REDUCING AGENTS FOR NITRATE Miscellaneous (Continued) GO. Murakami, T., "Volumetric Determination of Nitric Acid and Nitrate by Reduction with Stannous Chloride," Bunseki Kagaku 7, 766-71 (1558). 81. Pozsi-Escot, E., "The Determination of Nitric Nitrogen by Reduction with the Aid of Aluminum-Mercury," Compt. Rend., lU£, 1380 (1910). 82. Pozzi-Escot, E., "The Reduction of Nitrates to Ammonia and a New Method of Determining Nitrates," Ann Chim. Anal., lU, UU5-6 (1910). 83. Stammer, K., ["Reaction of HN03 and CO?" - Actual Title Unknown] Pogg. Ann. 82, 137 (1851). Cited in Gmelins Handbuch der anorganischen Chemie, 8th Edition, Syst. No. U, p. 1006 (1955). 8U. Thomas, M., "Total Determination of the Nitric Nitrogen and of the Nitrogen of Nitrated Groups by Titanium Chloride and Gravimetric Determination of a Nitrate Derivative in a Mixture with a Nitrate," Mem. Poudres 3Jt, 357-6? (1952). 85. Wood, E. D., Armstrong, F.A.J., and Richards, F. A., "Determination of Nitrate in Sea Water by Cadmium-Copper Reduction to Nitrite," J. Mar. Biol. Ass. U.K. V? (1), 23-31 (1967). D. DEAMMINATION AGENTS Sulfamic Acid 1. Baumgarten, P., "The Effect of Nitric Acid upon Sulfamic Acid. A Simple Method for the Preparation of Nitrous Oxide," Ber. JIB, 80-1 (3.938). 2. Baumgarten, P., and Marggraff, I., "The Reaction of Nitrites with Amino sulfonic Acids and the Detection and Estimation of Nitrous Acid in the Presence of Nitric Acid," Ber. 63, 1019-2U (1930). 3. Brasted, R. C., "Detection of Nitrite and Sulfamate Ions in Qualitative Analysis," J. Chem. Education 28, 592-3 (1951). U. Brasted, R. C., "Reaction of Sodium Nitrite and Sulfamic Acid," Anal. Chem. 2k, 1111-lU (1952). 5. Carson, W. N. Jr., "Gasometric Determination of Nitrite and Sulfamate," Anal. Chem. 23, 1016-19 (1951). 6. Wu, Ching-Hsien, and Hepler, L. G., "Thermochemistry of Sulfamic Acid and Aqueous Sulfamate Ion," J. Chem. Eng. Data 7, Pt. 1, 536-7 (1962). 59 ------- D. DEAMMINATION AGENTS Sulfamic Acid (Continued) 7. Gumming, W. M., and Alexander, W. A., "Use of Aminosulfonic Acid in the Determination of Nitrites," Analyst 68, 273-1* (19^3). 8. Groh, H. J. Jr., and Russell, E. R., "Intermediates Formed in the Reaction of Nitrite with Salts of Sulfamic Acid," J. Inorg. Nucl. Chem. 26 (1*). 665-7 (19610. — 9. Gottfried, J., and Novak, Jiri V. A., "Polarographic Determination of Amidosulfonic Acid and Nitrate," Chem. Prunysl 7, 1*76-8 (1957). 10. Heubel, J., and Canis, C., "Reaction Between Nitrates and Sulfamates," Compt. Rend. 255. 708-10 (1962). 11. Heubel, J., and Wartel, M., "The Reaction Between Nitrites and Amino Sulfonates," Compt. Rend. 257 (3), 68U-6 (1963). 12. Kaloumenos, H. W., "Specific Determination of Nitrite," Werkstoffe u. Korrosion n, 626 (1960). 13. Kostrikin, Yu M., and Mikhailova, N. M., "Treatment of Heating Water " USSR 139,997, Appl. Oct. 8, 1960. 1U. Subrahmanyan, P.V.R., Sastry, C. A., and Pillar, S. C., "Determination of the Permanganate Value for Waters and Sewage Effluents Containing Nitrite," Analyst 8J*, 731-5 (1959). Urea (See Also References D-lU,E-5) 15. Asendorf, E., "Method of Decontaminating Aqueous Solutions of Nitrites of Alkali Metals and/or Alkaline Earth Metals," B.P. 1,028,161 (Assigned to Water Engineering Limited), 1* May 1966. 16. Bonner, W. D., and Bishop, E. S., "The Hate of Reaction of Nitrous Acid and Urea in Dilute Solutions," J. Ind. Eng. Chem., £, 13^-6 (1913). 17. Burriel, F., and Suarez Acosta, K., "Analytical Problem of Separating Nitrates and Nitrites. IV. Destruction of Nitrites with Urea and its Derivatives," Anales Heal Soc. Espan. Fis. y. Quim. U6_B, 1*29-1*0 (1950). 18. Gorenbein, E. Ya, and Sukhan, V. V., "Interaction of Urea with Nitric Acid in an Aqueous Medium," Zh. Neorgan. Khim. 10 (7), 1701-5 (1965). 19. Sabbe, W. E., and Reed, L. W., "Investigation Concerning Nitrogen Loss Through Chemical Reactions Involving Urea and Nitrite," Soil Sci. Soc Am. Proc. 28 (1*), 1*78-81 (196!*). 60 ------- D. DEAMMINATION AGENTS Urea (Continued) 20. Shaw, W.H.R., and Bordeaux, J. J., "The Decomposition of Urea in Aqueous Media," J. Am. Chem. Soc. 77, u729-33 (1955). 21. Weston, C. F., "Removing Nitrous Acid from Solutions Such as Those of Sodium Nitrate," U.S. 2,139,11*2, Dec. 6, 1939. Amino Acids 22. Austin, A. T., "Deammination of Amino Acids by Nitrous Acid with Par- ticular Reference to Glycine. The Chemistry Underlying the Van Slyke Determination of a-Amino Acids," J. Chem. Soc. 1950, 1U9-57. 23. Cristol, P., Benezech, C., and Lissitsky, S., "Deammination by Nitrous Acid. I. Rate Constant of Deammination of Amino Acids in Aqueous Solution," Bull. Soc. Chim. Biol. 31, 150-6 (19^9). 2k. Cristol, P., Benezech, C., and Lissitsky, S., "Deammination by Nitrous Acid. II. Influence of Iodine on the Rate of Deammination of Amino Acids. Deammination of Mixtures of Amino Acids," Bull. Soc. Chim. Biol. 31, 156-60 (19^9). Miscellaneous (Ammonia, Amines, Azides) 25. Adamson, D. W., and Kenner, J., "Decomposition of the Nitrites of Some Primary Aliphatic Amines," J. Chem. Soc. 193U, 838-M*. 26. Burriel, F., and Saurez, R., "Analytical Problem of Separating Nitrates and Nitrites. III. Destruction of Nitrites with Ammonium Salts," Anales Real Soc. Espan. Fis. y Quim. U5B, 893-910 (19U9). 27. Kezdy, F. J., Jaz, J., and Bruylants, A., "The Kinetics of the Effect of Nitrous Acid on Amides. I. General Method," Bull. Soc. Chim. Beiges 6_7, 687-706 (1958). 28. Huckel, W., and Wilip, E., "The Conversion of Amines with Nitrous Acid," J. Prakt. Chem. 158. 21-32 (19U1). 29. Mohrig, J. R., "The Synthesis and Nitrous Acid Deammination of Some Bicyclic Amines. The Mechanism of the Deammination Reaction," Univ. Microfilms, Order No. 6U-U369; Dissertation Abstr. 25 (2), 8U2 (196^). 30. Ridd, J. H., "Nitrosation, Diazotization and Deammination," Quart. Rev. (London) 15_ (U), Ul8-»H (1961). 31. Seel, F., Wolfle, R., and Zwarg, G., "Kinetics of the Decomposition of Nitrous Acid with Hydrazoic Acid," Z. Naturforsch. 13b, 136-7 (1958). 61 ------- D. DEAMMINATION AGENTS Miscellaneous (Continued) 32. Spence, L. U., Whitmore, F. C., and Suraatis, J. D., "Action of Methyl- amine with Nitrous Acid," J. Am. Chem. Soc. 63, 1771 (19!*!). 33. Stedman, G., "Mechanism of the Azide-Nitrite Reaction. Pt. 1." J. Chem. Soc. 1959. 29U3-9- 31*. Streitwieser, A. Jr., "Reaction of Aliphatic Primary Amines with Nitrous Acid," J. Org. Chem. 22, 86l-9 (1957). E. CATALYSIS (See Also Individual Entries in Section D; e.g. D-20, -21) 1. Azim, M. A., and Saraf, S. D., "Catalytic Decomposition of Nitrous Acid." J. Indian Chem. Soc. 33, 763-1* (1956). 2. Azim, M. A., and Shafi, M., "Kinetics of Nitric Acid Decomposition in Liquid Phase," J. Nat. Sci. Math. 5 (2), 223-6 (1965). 3. Catalina, L., "Vanadium and Reduction of Nitrates in Plants. III. Acti- vation of the Nitrates in Presence of Vanadium," An. Edafol. Agrobiol. (Madrid) £2 (11-12), 731-5 (1966). U. Halpem, J., "The Catalytic Activation of Hydrogen in Homogeneous, Heterogeneous and Biological Systems," Advances in Catalysis. Volume XI, Academic Press, Inc., New York, 1959, pp. 309-11. 5. Quartaroli, A., "The Kinetics of Febrile Reactions. Contribution to the Study of Autocatalysis," Gazz Chim. Ital. 53, 31*5-68 (1923). 6. Suzawa, T., "Decomposition of Nitrous Acid in Aqueous Solution," Kagaku to Kogyo (Osaka) 31, 55-60 (1957). F. ANALYTICAL METHODS 1. American Public Health Association, American Water Works Association and Water Pollution Control Federation, Standard Methods for the Examination of Water and Waste Water , Publication Office, American Public Health Association, New York, N.Y., 1965 (12th edition), pp. 166-208. 2. Armstrong, F.A.J., "Determination of Nitrate in Water by Ultraviolet Spectrophotometry," Anal. Chem. 35, 1292-U (1963). 62 ------- F. ANALYTICAL METHODS (Continued) 3. Bastion, R., Weberling, R., and. Palilla, P., "Ultraviolet Spectro- photometric Determination of Nitrate," Anal. Chem. 2£, 1795-7 (1957). U. Dukes, E. K., and Wanace, R. M., "Stability of Ferrous Sulfamate in Nitric Acid Solutions," Contract No. AT/07-2/1, E. I. Dupont DeNemours and Co., Savannah, Feb. 5. Feigl, F., Spot Tests in Inorganic Analysis, Elsevier Publishing Co., New York, N.Y., 1956, pp. 32$ »7 6. Fisher, F. L., Ibert, E. R., and Beckman, H. F., "Inorganic Nitrate, Nitrite, or Nitrate-Nitrite," Anal. Chem. 30, 1972-U (1958). 7. Snell, F. D., and Snell, C. T., Colorimetric Methods of Analysis. IV. D. Van Nostrand Company, Inc., Princeton, N. J., 195^, pp. 317. 8. Walton, H. F., Principles and Methods of Chemical Analysis. 2nd Edition, Prentiss-Hall, Inc., Englewood, N. J., 1964, pp. 327-28. 9. Watt, G. W., and Chrisp, J. D., "Spectrophotometric Method for the Determination of Urea," Anal. Chem. 26, 1*52-3 (195*0 . 63 ------- APPENDIX ANALYTICAL PROCEDURE FOR THE SIMULTANEOUS DETERMINATION OF NITRATE AND NITRITE IONS Nitrate and nitrite can be simultaneously determined from a solution con- taining insoluble matter, either with or without a deammination agent present, using ultraviolet spectroscopic methods. Apparatus and Reagents Gary lU recording spectrophotometer with a set of matched 1 cm silica cells. Standard volumetric glassware. 70-72$ Perchloric acidj Baker "analyzed" grade. Sulfamic acid; 99^ Eastman white label. Sodium nitrate; Baker "analyzed" grade. Sodium nitrite; Baker "analysed" grade. Sodium nitrate stock solution; prepare by dissolving 0.6071 gm of sodium nitrate in 1 liter of distilled water to form a stock solution of 100 ppm nitrogen. Sodium nitrite stock solution; prepare by dissolving 0.^929 gra °f sodium nitrite in 1 liter of distilled water to form a stock solution of 100 ppm nitrogen. Preparation of Calibration Curves Make up a set of six standard solutions of mixtures of sodium nitrate and sodium nitrite in distilled water such that each solution contains a total of 50 ptm nitrogen. Also make up a blank. Prepare these solutions by diluting the following volumes of stock nitrate and nitrite solutions to 50 ml: 1. 0.00 ml N03 25.00 ml NOg 2. 5.00 ml NO" 20.00 ml NOg 3. 10.00 ml NO" 15-00 ml N0~ ------- U. 20.00 ml NO" 10.00 ml NO" 5. 20.00 ml NO" 5.00 ml NO" 6. 25.00 ml NO" o.OO ml NOl •5 2 The blank should contain only distilled water. The concentrations of the above solutions are as follows: 1. 0 ppm NO~-N 50 ppm NO"-N 2. 10 ppm NO~-N kO ppm NO"-N 3. 20 ppm NO^-N 30 ppm NOg-N ^. 30 ppm NO~-N 20 ppm NOg-N 5. UO ppm NO"-N 10 ppm NO~-N 6. 50 ppm NO~-N o ppm NO~-N The blank is zero in both. ™?\0f each standard into clea* 100 ml volumetric flasks and -90 ml of distilled water. Do not dilute to the mark. Check the matched silica cells for cleanliness by scanning from 250-1Q.5 Sii^SrV^J distin:d water in b°th cel^. If the baseline varies more than 0.005 absorbance units, reclean the cells. + *v,T° th! fi£!* s°J;ution» add k drops of 70-72$ perchloric acid and dilute to the mark. This should lower the pH to 2. Immediately scan from 2^0-igl millimicrons using distilled water as a reference. Read at 200 millimicrons and call this absorbance A. Add a threefold excess (approximately U mg) of ^i^nf11^0 aCld t0 the flask" m* wen "4 scan this solution from 250-195 millimicrons. Read this at 200 millimicrons and call this absorbance B. Repeat for each standard solution and the blank. soroance +>, + Si^.fulfamic acid has a slight absorbance at 200 millimicrons (1/500 that of N03) care must be taken not to add too large an excess and to add a iri^COnStant am°Unt to both san^le "d blank' Correct the absorbances A and B by subtracting the appropriate blanks. -wo-ncea A Absorbance B is proportional to the N0§ concentration. Plot ppm nitrate nitrogen vs absorbance B. Can this plot curve 1. ™.w«e Absorbance A minus absorbance B is proportional to N0£ concentration. Plot ppm nitrite nitrogen vs absorbance A minus absorbance B. Call this plot curve 2. 65 ------- Procedure with Deajamination Agent Absent If the sample has insoluble matter, filter about 15 ml through Whatman k2 filter paper. Follow the procedure outlined in the calibration curve discussion with the following variations. If the sample started out at 50 ppm NO§-N, dilute 5 ml to 25 ml. A reagent blank is essential and must be treated exactly like a sample. The reagent blank should include the reducing agent, catalysts, be the identical pH of the sample, and treated under the same temperature and time conditions of the sample. Calculations Read all values at 200 millimicrons. To get the true value for each absorbance, the appropriate blank must be subtracted before any calculations. Absorbance B is absorbance due to NO"-N. Calculate ppm nitrate-N from Curve 1. •* Absorbance A minus absorbance B is absorbance due to nitrite-N. Calcu- late ppm nitrite-N from Curve 2. Multiply by the dilution factor to obtain the original nitrate and nitrite concentrations. Interferences Most anions interfere somewhat at 200 millimicrons. However, they are much weaker absorbers than either N0§ or N0£. If the reagent blank has been carefully made up and handled, these interferences can easily be subtracted out giving rise to very little error. ALTERNATE PROCEDURE FOR SOLUTIONS CONTAINING A DEAMMHiATION AGENT Calibration A third curve must be prepared for nitrite ion at neutral pH's. Scan a series of nitrite solutions containing 0-2 ppm nitrite nitrogen in distilled water, from 250-195 millimicrons. Read the absorbance values at 200 millimicrons. Plot the absorbance at 200 millimicrons vs ppm NOl-N. Can this Curve 3. Procedure Two aliquots of each solution must be taken. The first aliquot is adjusted to pH between 7 and 9, diluted to the mark, and scanned "as is" 66 ------- from 250-195 millimicrons. Read the absorbance at 200 millimicrons and call - this value absorbance C. Treat the second aliquot exactly as before to obtain absorbances A and B. Absorbances A and B are equal if the deammination agent present is there in sufficient quantity to deamminate the solution as the pH is lowered to 2. Calculations Absorbance B is due to the absorbance of NO" ion. Read from Curve 1. Absorbance C minus absorbance B is due to the absorbance of NOZ ion. Read from Curve 3. 2 Calculate the results in the same manner as shown in the preceding section. 67 ------- |