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
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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
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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
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, 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
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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
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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.
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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
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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
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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.
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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"
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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.
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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.
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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
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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)
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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.
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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
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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
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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. Special Publication No. 10, The Chemical Society,
London, 1957, pp. 131-136.
U. Gehlen, H., "Reactions and Properties of Nitric Oxide and its Compounds.
II. The Salts of the Nitric Oxide Compound of Sulfurous Acid," Ber.
6_5_B, 1130-Uo (1932).
Occurrence and Effects
5. Azad, H. S., and King, D. L., "Effect of Industrial Waters on Lagoon
Biota," Purdue Univ., Eng. Bull., Ext. Ser. No. 118, UlO-22 (1965).
6. Biffoli, R. , "Determination of the Nitrate Ion and its Hygienic Impor-
tance in Tap Water. II. Results of Recent Controls in the Province of
Florence." Bon. Lab. Chim. Provincial! (Bologna) 16 (5), 558-70 (1965).
7. Burden, E.A.W.J., "The Toxicology of Nitrates and Nitrites with Parti-
cular Reference to the Potability of Water Supplies - a Review,"
Analyst 86, ^29.33 (1961).
8. Darvas, I., "On the Correlation Between the Nitrate Content of Bacterio-
logically Contaminated Waters and the Incidence of Methemoglobinemia,"
Egeszsegtudomeny 10 (4), 357-66 (1966).
9. DeMarco, J., Kurbiel, J., Symons, J. M., and Robecka, G., "Influence of
Environmental Factors on the Nitrogen Cycle in Water," J. Amer. Water
Works Ass. 5_£ (5), 580-92 (1967).
10. Fair, G. M., "Protecting the Purity of Inland Waters," J. Sanit. Eng.
Div., Am. Soc. Civil Engrs. go (SA6) 1-11 (196U).
11. Fassett, D. W., "Nitrates and Nitrites," Nat. Acad. Sci. - Nat. Res.
Counc., Publi. No. 1351*, 250-6 (1966).
-------
A. NITRATES IN WATER
Occurrence and Effects (Continued)
12. Ferguson, F. A., "A Nonmyopic Approach to the Problem of Excess Algal
Growths," Envir. Sci. Technol. 2, 188-193 (1968).
13. Feth, J. H., "Nitrogen Compounds in Natural Water - a Review " Water
Resources Res. 2 (l), 1+1.58 (1966).
1U. Flaigg N. G., and Reid, G. W., "Effects of Nitrogeneous Compounds on
Stream Conditions," Sew. and Ind. Wastes 26 (9), ll^-jU (195?).
15' frr™n«C;-R'J TWutrieS! Budfiet: Rational Analysis of Eutrophication in
a Connecticut Lake," Envir. Sci. Technol. 1 (5), 1*25-8 (1967).
16. g^^-J^*^ ™ Salt and Potable Water," Kiel Meeresforsch.
17. Hanson, A. M., and Flynn, T. F., "Nitrogen Compounds in Sewage," Purdue
18. Hollis, M. D., "The Water Ponution Situation - Aspirations and Realities "
J. Water Pollution Control Federation 37, 1-7 (1965). -""es,
19. Imhoff, K., "Nitrates Again. A Summary, " Gesundh.-Ing. 6Jt, 632 (19!*!).
*°* JSSJ1; «; i" 5"? "Mfl^r, S. A., "Excessive Nitrite Nitrogen in
Sludee> J' Wafcer *>^tion control Federation 33, 1286-9
21. Jtopanadze, Sh. Kh "Maximum Permissible Concentration of Nitrates in
Water," Gigiena i Sanit. 26, No. 9, 7-11 (1961).
.- ln "^ water>"
23' SSfitt: s'(8tei3?"J (^r*" on Mater Use>" J" *"• wster
2U. McCarty, P. L., et al, "Nutrient^Associated Problems in Water Quality
and Treatment, J. Amer. Water Works Ass. 5J3 (10), 1337-55 (1966).
25. Nichiporovich, A. A., "Photosynthesis and Mineral Fertilizers "
Agrokhimiya. 196!* (l), Uo-52. '
26. O'Brien J E., and Rosenthal, B. L., "Sman Activated Sludge Plants
Nitrate-Alkalinity-pH Relationship," Sanitalk 9_ (l), 18-21(1960-1) ? '
-------
A. NITRATES IK WATER
Occurrence and Effects (Continued)
27. Petr, B., and Schmidt, P., "The Influence of Altered Living Environment
on Children. II. Erythrogram and Methemoglobinemia in Blood," Ccsk.
Pediat. 21 (6), 505-5 (1966).
28. Kueffer, H., "Nitrification and Denitrification in Waste Water Purifi-
cation," Vom Wasser 31, 13^-52 (196U).
29. Sturm, G., and Bibo, F. J., "Nitrate Content of Drinking Water, Especi-
ally in the Rheingau Region," Gas-Wasserfach 106 (12), 332-U (1965).
30. U. S. Department of Health, Education and Welfare, Public Health Service
Drinking Water Standards, Public Health Service Publication No. 956,
U. S. Government Printing Office, Washington, D. C., 1962, pp. U7-51.
Natural and Biological Denitrification (See also Ref. A-55)
31. Allison, F. E., "Losses of Gaseous Nitrogen from Soils by Chemical
Mechanism Involving Nitrous Acid and Nitrites," Soil Sci. 9_6 (6),
1*6^-9 (1963).
32. Brandon, T. W., and Grindley, J., "Effect of Nitrates on the Rising of
Sludge in Sedimentation Basins," Surveyor 10J*, 7-8 (19^5).
33. Brezonik, P. L., and Lee, G. F., "Sources of Elemental Nitrogen in
Fermentation Gases," Air and Wat. Pollut. Int. J. 10, 1^5-160 (1966).
3U. Camp, T. R., Water and Its Impurities. Reinhold Publishing Corporation,
New York, 1963, pp. 280-2.
35. Greenwood, D. J., "Nitrification and Nitrate Dissimulation in Soil,"
Plant and Soil XVII (3), 365-77 (1962).
36. Johnson, W. K., and Schroepfer, G. J., "Nitrogen Removal by Nitrifi-
cation and Denitrification," J. Water Pollution Control Federation _3§,
1015-36 (196*0.
37. Jones, D. I. H., and Griffith, G. ap, "Reduction of Nitrate to Nitrite
in Moist Feeds," J. Sci. Food Agr. 16 (12), 721-5 (1965).
38. 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. Hsia, Shu-Fang, "Dark Reduction of Nitrate by Wheat Leaves," Sheng Wu
Hua Hsueh Yu Sheng Wu Wa Li Hsuch Pao 2 (2), 131-3 (1962).
^5,. Silver, W. S., "Enzymatic and Non-Enzymatic Reactions of Nitrate in
Autotrophic and Heterotrophic Microorganisms," 7th Intern. Congress
of Soil Science (Madison, Wis.), 592-9 (1960).
U6. Robert A. Taft Sanitary Engineering Center, Summary Report. Advanced
Waste Treatment Research Program. July. 196U - July. 1967. Publication
WP-20-AWTR-19, U. S. Department of the Interior, Federal Water Pollution
Control Administration, Cincinnati, Ohio, 1968, pp. 65, 68-69.
U7. Thames Survey Committee and Water Pollution Research Laboratory, Water
Pollution Research Paper No. 11, Effects of Polluting Discharges on the
Thames Estuary. Her Majesty's Stationery Office, London. 19^
pp. 247-255, 537.
U8. Vialard-Goudou, A., and Richard, C., "Reduction by Iron, of Nitrate Ions
to Nitrite and Ammonium Ions in Acidic Tropical Waters," Compt. rend
21+1, 978-80 (1955).
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
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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
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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
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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
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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
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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
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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
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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
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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
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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
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