Technical Bulletin No. 96
September 15, 1942
of Arizona
Sidney P. Osborn (ex officio), .
—.—................Governor of Arizona
E. D. King, BA. (ex officio).. State Superintendent of Public Instruction
Albert M, Crawford, B.S., President.,,
William H. Westover, U-.B.........
Martin Gentry, LL.B....
Cleon T. Knapp, LLJB., Treasurer
Jack B. Martin, Secretary.....
M. O. Best
Clarence E. Houston, 1X.B., B.A
Mrs. Joseph Madison Greer, B.A
Alfred Atkinson, D.Sc
President of the University
Paul S. Burgess, PhD
Ralph S. Hawkins, PhD.
William T. McGeorge, M.S
Agricultural Chemist
James F, Breazeale, B.S
Theophil F. Buehrer, Ph,D
Physical Chemist
Howard V, Smith, M.S......................... .........Associate Agricultural Chemist
William P. Martin, PhD
Assistant Soil Microbiologist
George E, Draper, M.S
...........Assistant Agricultural Chemist (Phoenix)
Alfred B. Caster, PhD,
Assistant Agricultural Chemist
Ammonia as Compared with Nitrate Nitrogen in Plant Nutrition....
Influence of Soil Alkalinity on the Nitrification Process
Toxic Effect of Ammonia on Nitrifying Bacteria
Description of Soils
Rate of Nitrification of Ammonia and Other Nitrogen Compounds
in Desert Soils
Influence of Initial pH Value on the Rate of Nitrite Accumulation
and the Formation of Nitrates
Series A—Control; No Treatment
Series B—Treated with NH t OH-fCa(OH) 2
Series C—NELOH Alone
Series D—Treatment with NHiOH+0.5tf H»SOt
Series E—Treatment with NHiOH+ltf H2SO,
Relationship between pH Value and Nitrate Formation
Correlation of Initial pH Value with Nitrite and Nitrate Formation 495
Effect of Maintaining the Soil at a pH Value above the Threshold
of 7.7±0.1
Energy Relations Involved in the Hydroxylamine-Hyponitrous
Acid Mechanism in the Nitrification of Ammonia
In recent years ammonia gas dissolved in irrigation water has
become agriculturally important as a source of nitrogen in the
fertilization of soils, particularly those of irrigated regions.
Gaseous ammonia, as distinguished from the common nitrogen
fertilizers, has several noteworthy advantages: (1) an unusually
high percentage of nitrogen—the highest, in fact, of all nitrogen
fertilizers; (2) the fact that by being wholly utilized by the plant
it leaves no salt residue to contribute to the salinity of the soil;
(3) the ease with which it can be applied to the soil even after
the crop is too high to permit the application of granular fertilizers by the conventional methods; and (4) the economy with
which the amount applied and its distribution can be controlled.
In addition, the loss of ammonia by volatilization when so applied
is remarkably slight on account of its prompt fixation by base
Superficially it may seem anomalous to apply a highly alkaline
ammonia solution as a fertilizer to an already alkaline soil, and
the established practice has generally been to use a nitrogen
compound having an acid or neutral reaction. Nevertheless, on
many soils where ammonia has been employed, it has produced
increases in yield of magnitudes such that the anticipated undesirable effects of high alkalinity were entirely overcome by the
rapidity with which the ammonia was oxidized. Since the rate of
this oxidation process is without doubt a factor affecting the crop
response which may be expected as a result of ammonia fertilization, it was deemed of interest to investigate the interrelationships
of the physical, chemical, and microbiological factors which influence it.
The authors are well aware that the mechanism of ammonia
oxidation in the soil has for many years engaged the attention of
investigators, and that its quantitative relationships are fairly
well known. In the present instance, however, it is evident that
there may be conditions that are unfavorable to the process,
either (1) by slowing down certain steps, (2) by allowing intermediate products toxic to the nitrifying bacteria to accumulate,
or (3) by inhibiting nitrification entirely. The extent to which
*The authors desire to express their appreciation to the Shell Chemical
Company of San Francisco, California, for sponsoring the fellowship under
which this investigation was carried out and for financing in part its
tShell Fellow, 1939-41.
these effects manifest themselves in any given soil is evidently a
function of various factors, notably its alkalinity. This bulletin
sets forth the results of experiments to determine to what extent
the alkalinity affects the process, and whether or not there is a
threshold pH value above which the conversion of ammonia to
nitrite, or of nitrite to nitrate, will not occur regardless of how
favorable the other conditions may be,
Inasmuch as the use of free ammonia as a fertilizer is a comparatively recent development in agricultural practice, the amount of
the published research on its properties and behavior is necessarily limited in extent. Furthermore it is not possible to deduce
with any degree of certainty the probable behavior of ammonia
when so used, from the results of investigations on the various
ammonium salts or of the organic nitrogen compounds commonly
applied as fertilizers. On the other hand, the background of existing research on the varied factors affecting the microbiological
oxidation of nitrogen in both the organic and inorganic forms and
the utilization of those forms by the plant has yielded valuable
suggestions as to an advantageous approach to the present
The ability of plants throughout their entire growth period to
utilize ammonia per se has not yet been conclusively proved.
Physiological differences in plants with respect to the intake of
other nutrient elements, the development of their root systems,
and fruiting characteristics have been observed repeatedly where
plants were grown in nutrient solutions or sand cultures in which
either the nitrate or ammonium ion was present as the sole source
of nitrogen. Some investigators, notably Prianishnikov (46),
claim to have established the fact that plants under certain conditions absorb ammonia in preference to nitrate. Studies on
ammonium-ion absorption have, however, generally involved the
use of neutral ammonium salts in nutrient solutions, in sand
cultures or in waterlogged soils—all of which represent conditions which do not obtain in normal, arable soils.
Shive and his co-workers (2, 14, 15, 20, 48, 49, 50), Jones and
Skinner (29), Stewart et al. (52), McGeorge (34), and Naftel (41)
have made substantial contributions to our knowledge of the
relative absorption and assimilation of ammonium ion and nitrate
by plants. Considered collectively, their results indicate that the
ion absorbed and assimilated is a function of the age and species
of the plant and/or the reaction of the external medium. The
preponderance of evidence, however, indicates that nitrates are
essential to the plant at some stage of its development and
growth. As noted recently by Arnon (1), the possibility of
nitrification of ammonium ion in any study which aims to compare its relative absorption with that of nitrate is a complicating
factor which renders the use of, or even the presence of, soil undesirable. On the other hand, since the conditions are not comparable, it is poor logic to assume that the results of nutrient
solution studies will apply directly to the soil. From an experimental point of view it would seem preferable to make a study
of the state or condition of the nitrogen in the soil at various intervals of time after its application—that is, to determine its nitrification rate—when subjected to conditions such as those which
actually prevail in the soil.
Although numerous compounds have been studied as a source
of nitrogen in fertilizer programs, practically no work bearing
either on field or laboratory studies with gaseous ammonia has
appeared in the literature. Ammonium sulfate has been extensively used as a standard for the comparison of different nitrogen
compounds in nitrification studies, but its residually acid nature
makes it difficult to apply the results so obtained in the prognosis
of the probable behavior of ammonia when applied to a soil. The
majority of the published papers concern themselves with nitrification studies on humid soils which are dominantly acid and
may possess a microflora in general different from that of the
dominantly alkaline soils of the desert.
Because of the alkalinity produced when ammonia is dissolved
in water, the pH factor so introduced may influence not only the
rate at which the ammonia is nitrified but also certain physical
properties of the soil. When ammonia is added in an amount
equivalent to 300 p.p.m. of nitrogen, it is sufficient not only to
neutralize the carbon dioxide and bicarbonate ion which the irrigation water may have contained but also to leave a certain
amount of free alkalinity. The effect of the added base on the pH
value of the soil will of course be determined by the extent to
which the latter is buffered against base, which in alkaline soils
is usually slight. In this manner an amount of free alkalinity is
added to the soil which, though temporary in duration, may nevertheless have important effects on the microbial population of the
soil, particularly the nitrifying organisms. Hence their ability to
remain active over the range of pH values which may be produced under such conditions may be in considerable measure reduced. These are some of the fundamental aspects of the problem
which received particular consideration in the course of the present study.
That nitrification proceeds more efficiently under neutral or
slightly alkaline conditions has been noted by various investigators in this field, but no study appears to have been made on the
optimum or limiting pH values for nitrification in alkaline calcareous soils. Fraps and Sterges (21) found that soils which
failed to nitrify ammonium sulfate could be made to do so by the
addition of cultures from actively nitrifying soils and/or of
calcium carbonate. These authors (22), furthermore, found that
in acid soils calcium carbonate favored rapid nitrification more
so than did dicalcium phosphate, rock phosphate, magnesium carbonate, or dolomite. Similar studies by Tandon and Dhar (53)
showed that nitrification is favored more by calcium carbonate
than by magnesium carbonate.
Waksman (54) in his nitrification studies stressed the necessity
of neutralizing the acid formed in the oxidation of ammonium
sulfate and indicated that an acid condition was harmful to the
process. He observed, for example, that the accumulation of nitrates stopped when a pH value in the range of 4.4 to 4.8 was
reached and concluded that nitrate accumulation in any soil depends upon its initial pH value, its buffer capacity, and the final
pH value of the soil, "more so than on the bacteriological activity."
The minimum pH value reported by Waksman is essentially in
agreement with that found by Humf'eld and Erdman (28) and by
Gaarder and Hagem (23); the latter, having worked with acid
solutions, placed the minimum pH range below which the nitrifiers no longer function between 3.9 and 4.5.
Naftel (42) investigated the rate of nitrification in five widely
different soils as a function of pH value and base saturation and
found, in general, that the rate increased with base and calcium
saturation. His results indicated that the extent of nitrification
in different soils may vary widely even though their pH values
may be the same. This may be due in part to the degree of saturation with bases, since only those soils which had the highest percentage of their base exchange complex saturated with calcium
nitrified ammonium sulfate to any considerable extent.
That the nitrifying bacteria are capable of adjusting themselves
to a fairly wide variation in pH value was found by Olsen (44)
working with strongly acid humus soils in which the pH value
had been adjusted and maintained at the desired levels with lime.
He concluded that nitrification can take place between pH 3.7 and
8.8, the optimum being 8.3, so long as ammonia is not the limiting
Meyerhof (37, 38, 39), Gowda (25), and Meek and Lipman (36)
have reported results which pertain more specifically to nitrification in the alkaline range. In a detailed study of both types of
bacteria, Meyerhof found that ammonia was oxidized most
quickly between pH 8.5 and 8.8. The optimum pH for nitrite
oxidation was between 8.3 and 9.3. These values are in substantial agreement with those obtained by Gowda. Meek and Lipman,
on the other hand, found that both nitrite- and nitrate-forming
organisms from garden soil were alive and functioning at pH 13.0,
but the nitrifiers from acid peat were unable to produce nitrates
above a pH value of 9.5. They concluded that the nitrate-forming
bacteria are somewhat more resistant to alkalinity than are the
nitrite formers. It is to be noted, however, that the above investigators worked exclusively with nutrient solutions and not with
Ayers and Jenny (4) and Waynick (56) studied the nitrification of ammonia and ammonium sulf ate in relation to the physical
and chemical changes resulting from the application of these fertilizers. The former used one acid and one slightly alkaline soil in
their studies. Waynick worked with a series of soils ranging from
7.1 to 8.1 in their initial pH values, but since his determinations
were made on the soil suspension, his pH values are obviously
higher than those obtained on the same soil at field moisture contents (35).
Nelson (43) found that the toxicity of manganese was reduced
by additions of lime to the soil and concluded that the shift in reaction so produced rendered the conditions more favorable for
the action of the nitrifying organisms.
There is some evidence to indicate that free ammonia may have
a toxic or inhibiting effect on the nitrifying organisms. Willis
and Piland (57) found that the free ammonia formed from the
hydrolysis of diammonium phosphate decreased the rate of nitrification when this salt was used in pot culture experiments.
Such toxicity, strangely enough, was not observed in the case of
ammonium sulf ate, chloride, or nitrate; nor did the alkalinity of
the diammonium phosphate appear to be responsible for the injurious effect. They found further that calcium salts were able
to counteract the toxicity of the ammonia. It is therefore reasonable to expect that such toxicity would not be likely to occur in
calcareous soils, notwithstanding the alkalinity produced when
free ammonia is added as the source of nitrogen.
Similarly Waksman (54) reported that sufficient free ammonia
is evolved in the rapid decomposition of dried blood in alkaline
and poorly buffered soils to have an injurious effect upon the
activity of the nitrifying bacteria in the soil. He did not, however,
mention the antagonistic effect of calcium as reported by Willis
and Piland, and since his report referred specifically to alkaline
soils, it is felt by the present writers that further studies should
be conducted on this phase of the problem to clarify the apparent
In view of the increasing interest in the use of gaseous ammonia
as a fertilizer for citrus, truck crops, small grains, and hay which
are extensively grown under irrigation on the alkaline calcareous
soils of the desert, it seemed desirable to study some of the fundamental aspects of its behavior when applied to such soils. It is
recognized that the rate and mechanism of ammonia nitrification
in such soils may be a function of the alkalinity, salinity, soil
type, and dominant microflora, and perhaps other factors which
influence it to a greater or lesser degree. The results to be presented in this bulletin concern themselves primarily with the pH
relationships which are involved in the nitrification mechanism.
The investigation was planned as a laboratory incubation study
under controlled conditions so as to afford a comparison between
the rate of microbial oxidation of ammonia with that of other
types of nitrogen fertilizers commonly used—namely, an inorganic
ammonium salt, ammonium sulfate, and an organic nitrogen
compound, urea. Six typical but widely separated desert soils,
having textures varying from sands to clays, were chosen. Three
of these were in the virgin state and three were from areas under
cultivation. All of them except one were calcareous and decidedly
The soils used in this study were as follows:
1. Superstition sand, from near the University Farm on the
Yuma mesa; virgin, calcareous.
2. Gila sandy loam, from a ranch adjacent to the Cortaro Farms
at Marana, Arizona; virgin, calcareous.
3. Gila sandy loam, same general locality as (2) but from a
field in cotton on Cortaro Farms at Marana, Arizona; under cultivation, calcareous.
4. Pima clay loam, from a farm in the Gila River bottom near
Safford, Arizona; under cultivation, mildly calcareous.
5. Laveen loam, from the University of Arizona Farm at Mesa,
Arizona; under cultivation, calcareous.
6. Palos Verdes sandy loam, from north of the Rancho Palos
Verdes near Tucson, Arizona; virgin, noncalcareous.
The samples were taken within a period of about a month during the summer of 1939. After gentle rolling to break up the aggregates, and removal of foreign matter, the soils were screened
to 10-mesh, thoroughly mixed, and stored.
Prior to the nitrification studies, it was deemed of value to
make determinations of certain chemical constituents and physical properties of these soils which might have some bearing on
the nitrification process. The following were accordingly determined by the accepted methods: pH values by the Beckman pH
meter; soluble salts by conductivity of the 1:5 aqueous extract;
total carbonate by the official gasometric method of the A.O.A.C.
(3); nitrate by phenoldisulfonic acid; organic carbon by wet oxidation with chromic acid; and total and specific buffer capacities
on the 1:5 suspension according to the method of Pierre (45).
The data are assembled in Table 1.
The pH values were determined both on the 1:5 suspension and
on the soil at a moisture content of 70 per cent of its water-holding
capacity. The data given in the table illustrate the frequently
observed fact that the pH value of a soil at field moisture content
is considerably lower than that of the same soil in the 1:5 suspension, and in most of the above soils the difference amounted to as
much as 1.1 pH unit. None of these soils at field moisture con-
tent can be said to be excessively alkaline, being, in fact, fairly
close to neutrality, and one of them, the Palos Verdes sandy loam,
actually has a value on the acid side of neutrality.
With the exception of the Pima and Laveen loams, the soluble
salt contents of these soils are relatively low, of a magnitude
which is not great enough to affect seriously the activity of the
nitrifying organisms. The high salt content of the two soils referred to accounts in part for their relatively lower pH values.
The data for total organic carbon are consistent with the typically low percentage of organic matter in desert soils. All of the
soils except the Palos Verdes sandy loam contained appreciable
amounts of carbonate, particularly the Laveen loam, to which the
latter soil owes its unusually high buffer capacity. The specific
buffer capacity data afford a significant indication of the stability
of these soils toward a change in pH. When the specific buffer
capacity is plotted against the total carbonate content, the points
fall on a straight line within limits of experimental error, thus indicating that calcium carbonate is the principal buffering compound present. For this reason and also because the soils on the
alkaline side of neutrality are largely saturated with bases, they
are buffered primarily toward acid. Being so slightly buffered
toward base, as the buffer determinations have shown, it is not
surprising that a considerable rise in pH is realized when solutions of ammonia or of other free bases are added to the soil. In
the case of the Palos Verdes sandy loam, these conditions are
Since buffer capacity determinations are usually made at a
constant dilution of 1:5, it is conceivable that a soil may exhibit a considerably higher buffer capacity at field moisture
content. This was found to be the case in later experiments when
equal volumes of normal and half-normal sulfuric acid solution
were added to samples of Gila sandy loam at field moisture content, with no significant difference in pH value. These buffer
considerations have an important bearing upon how a given soil
will react when treated with ammoniated irrigation water and
the extent to which such a soil will resist any considerable decrease in pH value upon its subsequent nitrification.
Some of the soils contained appreciable amounts of nitrate and
organic matter, particularly those which have been under cultivation for some time. This is illustrated in the case of the Pima
clay loam, one of the most fertile agricultural soils of Arizona,
whose fertility can be accounted for by the fact that the farm
from which the sample came had been systematically fertilized
with manure at the rate of 12 tons per acre for the past 10 years.
The foregoing determinations are presented to emphasize the
fact that the soils chosen for the nitrification studies to be set
forth in this bulletin represent a fair cross section of the agricultural areas where ammonia is already being used, or may be used,
as a nitrogen fertilizer—a circumstance which must be borne in
mind in making generalizations concerning the microbiological
transformations it may undergo under field conditions.
The first series of experiments in this investigation was designed
to ascertain not only the general nature of the nitrification curve
for ammonia, ammonium sulfate, and urea over an extended period of incubation but also the nature of the process insofar as it
may be affected by the character of the soil.
Replicate 100-gram samples of soil were weighed into glass
tumblers with close-fitting covers. The samples were divided
into four equal groups and treated with the following fertilizer
solutions: Group 1: control, untreated; Group 2: 30 mg. of nitrogen as ammonia; Group 3: 30 mg. of nitrogen as ammonium
sulfate; Group 4: 30 mg. of nitrogen as urea. On this basis the
nitrogen was present in each sample to an extent of 300 p.p.m. of
air-dry soil.
Since ammonia is fixed very promptly by the soil, there is no
perceptible loss when so applied. By keeping the ammonia solution in a closed delivery system, its concentration could be maintained constant. Following application of the above solutions
distilled water was added to bring the moisture content of each
sample to 70 per cent of the water-holding capacity of the particular soil. The tumblers were then weighed and placed in an incubator maintained at 30 degrees C. The moisture lost during the
incubation was periodically restored by bringing the samples
back to their original weights by the addition of distilled water.
After the desired incubation intervals, duplicate samples from
each series were removed from the incubator and the pH values
determined on the moist soil. The entire 100-gram sample of soil
was then transferred to a large wide-mouthed bottle, 500 ml. of
distilled water added, and the resulting 1:5 suspension shaken
for 20 minutes to bring the nitrites and nitrates, formed during
the incubation, into solution. Carbon dioxide was then bubbled
through the suspension for a few minutes to coagulate the colloids
and facilitate the filtration. The clear filtrates were then analyzed
for nitrates by the standard methods, the final colorimetric determinations being made with a Cenco Photelometer. The data
are reported in terms of parts per million of nitrogen as nitrate
on the basis of the air-dry soil. Throughout this investigation the
analyses were made on duplicate samples of each treatment of
each soil. This procedure eliminates possible systematic errors
resulting from the use of aliquot portions of only one soil sample
and, in addition, the error arising from variation in the soil from
sample to sample. Thus each analytical result, as reported in the
tables and graphs to follow, is an average of determinations on
two independent soil samples and hence proportionately more
significant and trustworthy. By the same token the curves,
although continuous, do not represent a continuous change in any
single sample but rather the changes in similarly treated, replicate
The results presented in Table 2 give the successive amounts
of nitrate observed in the three respective treatments on the six
soils over a period of about 100 days.
The results for Laveen loam are shown in Figure 1. Because
of limitations of space and similarity in the general character of
the nitrification curves, it was deemed desirable to limit the
graphs to one of the soils in which the trends of the nitrification
process were fairly well exemplified.
It will be noted that the curves representing the rate of nitrification are typically exponential and hence are similar to the
growth rate curves for the nitrifying bacteria as found by numerous investigators (12). In the cultivated Gila sandy loam,
Laveen loam, and Pima clay loam, the rates of nitrification of
ammonia, ammonium sulfate, and urea were practically identical
with the exception that in the cultivated Gila sandy loam the
ammonia was oxidized somewhat more slowly than either of the
other two nitrogen compounds. It should be noted incidentally
that the corresponding curve for the control sample exhibits only
a very slight increase in nitrate during the first several days of
incubation and thereafter remains horizontal over the entire
period of incubation. This slight increase in nitrate probably resulted from the nitrification of the organic nitrogen originally
present in the soil.
The remaining three soils—namely, the Superstition sand, virgin Gila sandy loam, and Palos Verdes sandy loam—differed in
their rates of nitrification from those discussed above. Referring
to Table 2, it may be noted that ammonium sulfate was nitrified
less rapidly than the urea or ammonia in the Palos Verdes sandy
loam, which is diametrically opposite to the behavior observed
in the other soils. This behavior is very probably due to the fact
that during the oxidation of the ammonium sulfate the pH value of
the soil is lowered sufficiently to inhibit nitrification. At the end
of 40 days' incubation and with only one third of the ammonia
nitrogen oxidized, the pH value had already dropped to 4.65,
which incidentally is in close agreement with that of Waksman
(54) and Humfeld and Erdman (28)—namely, 4.4—as the "limiting reaction for nitrification on the acid side of neutrality.
In the Superstition sand (Table 2), nitrates did not begin to
accumulate for nearly 30 days, regardless of the nature of the
nitrogen fertilizer added, and even then such accumulation occurred only in the case of ammonium sulfate and urea. Not even
a trace of the ammonia was found to have been oxidized during
the entire 107 days of incubation. That the rate of nitrification
was rather slow in the case of Superstition sand is not surprising
considering its inherently low native fertility. It is most surprising, however, that none of the ammonia had been oxidized,
whereas the urea and ammonium sulfate were readily nitrified,
The data in Table 2 show that the ammonia solution had raised
280 —
Kd Treatment
300 p.p.m.(N) as NH4OH —
300 p.p.m.(N) as (NHA)PSOA
300 p.p.m.(N) as Urea
Figure 1.—Relative rates of nitrification in Laveen loam under different
the initial pH value of this soil to 9.5, and that the alkalinity had,
during the incubation period, decreased to a value equivalent to
pH 8.35 corresponding to the urea but higher than that of the
ammonium sulfate—namely, 8.21—which existed in the soil prior
to the beginning of the incubation. This observation, in addition
to the anomalous behavior of ammonia in virgin Gila sandy loam,
next cited, was the first definite indication that high alkalinity may limit, or even prevent, the oxidation of ammonia to nitrate
in calcareous desert soils.
In the ammonia-treated samples of virgin Gila sandy loam, no
nitrates were observed until the forty-ninth day and then only
in one of the duplicate samples. Since duplicates had in all of the
previous instances been in good agreement, it was at first thought
that this relatively enormous discrepancy was due to an error
which had crept into the technique employed in the nitrate determination on one of the samples. Particular care was therefore
taken with the corresponding samples taken off on the fifty-third
day, but again a considerable difference in nitrate content between the duplicate samples was noted. Throughout the remainder of the incubation, the behavior of the samples was unusual; on the fifty-ninth day nitrates were entirely absent in both
samples; on the seventy-second they were present in one but not
the other; on the eighty-sixth nitrates were present in considerable amount in both samples; and on the one hundred and seventh
the amount of nitrate present was again of the same order of
magnitude as in the control.
This peculiar behavior was difficult to account for, except insofar as it might be related to the pH values of the samples in
question. The data of Table 2, as well as the curves in Figure 2,
show, in fact, that there is a close correlation between the pH
value and the presence or absence of nitrates. It will be noted
from Figure 3 that on the forty-ninth day, When nitrates first
appeared in one of the samples, the pH value of this sample was
7.9. In all other instances the failure of the ammonia to be nitrified correlated closely with a pH value of 7.7 it 0.1. Similarly, all
of the samples which contained nitrates exhibited a pH value below this figure. As a consequence of this striking coincidence, the
data for all of the soils tested were re-examined, with the result
that in only one instance out of 133 in which nitrification was
observed to have taken place did nitrates make their appearance
at a pH value above 7.7 ±. 0.1.
In view of this finding and of the supporting evidence which is
to follow, the authors feel justified in concluding that: There
appears to exist in alkaline desert soils a threshold pH value of
7.7 zb 0.13 above which the complete oxidation of ammonia to
nitrate will not occur, and to which the pH value of such soils
must first be reduced before nitratification will take place.
The foregoing conclusion is of unusual interest and significance,
particularly in view of the results of Meek and Lipman (36) who
placed the optimum value at pH 8.5 to 8.8; Olsen (44) at 8.3, and
Meyerhof (37, 38, 39) between 8.3 and 9.2. The significance of
such a value in connection with the microbiological oxidation of
ammonia lies in the fact that in desert soils their inherent alkalinity must in some manner be reduced, and that the organisms are
virtually ineffective, so far as the primary nitrification process is
100 days
Figure 2.—Relative changes in pH value and nitrate content of virgin Gila
sandy loam treated with 300 p.p.m. of nitrogen as ammonia.
concerned, until the pH value of the soil has been reduced to the
threshold value.
The foregoing generalization regarding the existence of a
threshold pH value for nitratification suggested certain confirmatory experiments, the results of which will now be presented.
In the previous sections of this bulletin the process of nitrification has been referred to in its broad sense, including all of the
intermediate stages through which nitrogen may pass in the
course of its oxidation from ammonia to nitrate. In discussing
the subsequent phases of this investigation, it will be advantageous
to divide the process into two main steps: the first, which will
hereafter be referred to as nitritification, includes all of the possible intermediate steps in the oxidation of nitrogen from any of
the lower forms in which it may occur to nitrite; and the second,
to be termed nitratification, will signify the last stage of the process in which the nitrite is oxidized to nitrate.
From the drop in pH value of the soil prior to nitrification, as is
evident from the data in Table 2, it follows that the nitrifying
bacteria native to such soils do not function effectively until a
threshold pH value of 7.7 has been reached. This fact is unique
in view of the prevailing opinion that a decrease in pH occurs concurrently with the oxidation of the nitrogen to nitric acid. In the
present study the pH drop in the incubated samples was found to
occur prior to the detection of nitrates. It is self-evident that the
formation of sulfuric and nitric acid must, or should, result in
reducing to some extent the alkalinity of the soil. If, therefore, a
pronounced decrease in pH value occurred before nitrates began
to be formed, it seems reasonable to assume (in the case of the
urea, and ammonia-treated samples in which no sulfate radical is
present) that the reduction in pH must have come about through
the formation of compounds other than nitric or sulfuric acids.
The over-all oxidation of ammonia to nitrite is represented by the
following equation:
2 NH3 + 3 O2 + 2 OH" = 2 N(V + 4 H2O
in which the reduction in pH value results from the fact that
hydroxyl ions are used up in the reaction. It is possible, of course,
for some intermediate nitrogen compound other than nitrite to be
Assuming that the foregoing reaction actually accounted for the
observed changes in pH, it seemed likely that the nitrite ion
might be detectable in the soil in appreciable amount during the
course of the oxidation, especially in those samples in which a considerable difference existed between the initial pH value of the
soil and the threshold value. Under normal conditions and in
arable soils, however, nitrites are considered to be too transitory
to accumulate in significant amounts. It was considered of great
interest, therefore, to determine whether or not nitrite may, if
formed under the alkaline conditions of desert soils, actually accumulate in measurable amounts, and if so, whether the reduction
in pH value observed could be explained in terms of such accumulation.
Experiments were accordingly planned embodying the following
phases of the problem: (1) the influence of the initial pH value of
the soil upon the rate of nitrite accumulation and the formation
of nitrates; and (2) the effect of maintaining the pH value at a
level above the threshold value by the addition of a base, such as
calcium hydroxide. In order to permit a greater variation of other
factors, this phase of the study was confined to the cultivated Gila
sandy loam, a soil of relatively low salt content, medium buffer
capacity, and especially one which in previous experiments had
been found to permit the rapid oxidation of ammonia, as shown
in Table 2 and Figure 3.
100 days
Figure 3.—Relative changes in pH value and nitrate content of cultivated
Gila sandy loam treated with 300 p.p.m. of nitrogen as ammonia.
As in previous experiments, 100-gram samples of soil were
weighed into glass tumblers with close-fitting covers. These were
then divided into five groups and subjected to the following treatments: (A) control—no treatment; (B) 10 ml. of ammonia solution containing 30 mg, of nitrogen plus 10 ml. of saturated calcium
hydroxide; (C) 10 ml. of the same ammonia solution plus 10 ml.
distilled water; (D) 10 ml. of the same ammonia solution plus 10
ml. 0.5N sulfuric acid; and (E) 10 ml. of the same ammonia solution plus 10 ml. of 12V sulfuric acid. These treatments served to
fix the initial pH values at different levels with respect to the
threshold value. The volumes of solution, totaling 20 ml. in each
case were so chosen that the resulting moisture content of the
samples was 70 per cent of the water-holding capacity of this soil.
In order to insure a rapid and approximately uniform distribution
of the added liquids throughout the soil mass, each sample of soil
was transferred from the tumbler to a clean sheet of paper and
divided into approximate fourths. Five milliliters of the solution
to be used were then put into the empty tumbler and one fourth
portion of the soil sample added to it. A second 5 ml. portion of
solution was added, then a second quarter of the soil, and so on,
until all of the solution and soil for each tumbler had been added.
In this manner the initial pH value of the soil was quickly established, and, more important still, an even distribution of the
ammonia was achieved throughout the soil mass. The samples
were again incubated at 30 degrees C., and the moisture lost by
evaporation was restored at intervals. Periodically, duplicate
samples from each treatment were removed and analyzed for
nitrite and nitrate, and the pH values on the moist soil were
The results obtained in this study are presented in Table 3 as
well as in Figures 4 and 5, in which p.p.m. of nitrate formed, as
well as the pH values, are plotted as ordinates against incubation
25 days
Figure 4.—Correlation between pH value and nitrate formation in Gila
sandy loam under different treatments.
time as abscissae. The purpose of this study was considered
achieved as soon as the pH value had dropped to a point sufficiently low so that nitrification could proceed rapidly. For this
reason the data presented in the afore-mentioned figures were
made to extend through the twenty-fifth day only. A* final analysis was made, however, on the fifty-eighth day to determine the
maximum amount of nitrate that had formed. This determination
was made primarily to compare the results of the C treatment
with the data obtained in the previous experiments on the rate
of nitrification of ammonia, which had been carried out on the
same soil and under identical conditions of incubation. The behavior of the various series of treatments will now be discussed.
Series A—Control; no treatment
Eeferring to Figure 4, it will be noted that the pH values and
nitrate contents of the samples of the A series remained practically unchanged over a period of 25 days. The initial pH value
was at the threshold level of 7.7, and in the subsequent samples it
increased to as high as 7.9. Nitrates were present to the extent of
only 8 p.p.m. Nitrites were found to be present only as mere
traces throughout the incubation period.
Series B—Treated with NH,OH+Ca(OH)2
This series of samples, as a result of the addition of calcium
hydroxide, exhibited the highest initial pH value—namely 9.21.
It will be noted from Table 3 that the pH value gradually decreased during the first 8 days of incubation to about 8.4, remaining practically constant at that value up to the seventeenth day.
Thereafter a sudden drop in pH occurred, accompanied by the
appearance of nitrite, and decreased gradually to 7.60 at the end of
the experiment. The concentration of nitrite in the soil samples
continued to increase, however, reaching a maximum of 71 p.p.m.
on the twenty-first day just as the pH dropped to the threshold
value. Not until the twenty-fifth day, while the pH value was still
at the threshold level, did nitrates make their first appearance and
nitrites begin to disappear.
Series C—NH^OH alone
The data of this series are of particular interest since they confirm the results obtained under the same experimental conditions
as in the nitrification rate study already reported in Table 2 and
Figure 3.
In the present instance the samples had a relatively high initial
pH value—namely, 9.0—and the decrease during the first 11 days
of incubation was approximately 0.6 pH unit. Up to the fourteenth
day, as shown by the data in Table 3, no additional nitrate had
been formed, but a definite quantity of nitrite (22 p.p.m.) had
made its appearance. As the build-up of nitrites continued, the
pH value dropped somewhat, and eventually fell to a point below
the threshold value by the time the maximum amount of nitrite—
namely, 94 p.p.m.—had accumulated. This relatively high maximum for nitrite was reached on the twenty-first day of incubation,
but thereafter the amount decreased sharply, so that by the
twenty-fifth day only a trace of nitrite was detectable in the
samples. Almost immediately after the pH value had dropped
below the threshold value (which occurred on the seventeenth
day), nitrates began to form and thereafter accumulate.
From the above data it is evident that nitrites must have been
present in the soil for nearly 2 weeks, a fact which is rather
surprising, in view of the generally accepted fact that nitrites do
not exist as such to any appreciable extent in well-aerated soils,
at normal moisture contents, being oxidized to nitrate almost as
soon as they are formed.
It is of interest to compare the nitrate data for the C series of
samples with those obtained in the initial nitrification experiment illustrated in Table 2 and Figure 3. It will be recalled that
the same Gila sandy loam (cultivated) had been fertilized with
ammonia to exactly the same extent—namely, 30 rng. of nitrogen—
and incubated under identical conditions. Referring to Table 3, it
will be noted that the first appreciable increase in nitrate content
occurred on the nineteenth day of incubation, and on the twentyfifth day nitrates were found to be present to an extent of 190
p.p.m., which represented a nitrogen recovery of about 63 per
cent. In the previous experiment shown in Figure 3, the first
evidence of nitratification was noted on the eighteenth day of
incubation, and on the twenty-fifth day a maximum of 193 p.p.m.
of nitrates was observed. Here the recovery was about 64 per
cent of the total nitrogen applied as ammonia.
Inasmuch as 97 per cent of the nitrogen applied in the first
series had been recovered at the end of 7 weeks, it was decided
to allow a final set of the samples to incubate for approximately
the same period—namely, 58 days—to observe how closely the
final values recorded in Table 3 would be confirmed. It was found,
in fact, that in the C series 99 per cent of the nitrogen initially
added as ammonia had been oxidized to nitrate by the end of this
This phenomenal agreement between two entirely independent
incubation studies indicates how reproducible the results of such
studies on nitrogen transformations can be and constitutes a
confirmation of the threshold pH value in nitratification processes
as had been repeatedly observed in the course of the present
Series D—Treatment with NH4OH-f 0.5N HoSO* and Series E—Treatment
with NHiOH+lN H2SO*
In this experiment an attempt was made to adjust the initial pH
value to a point below the threshold, by adding equal volumes of
0.5JV and IN sulfuric acid to the 100-gram samples of soil after
treatment with ammonia. From the difference in acid concentration it was anticipated that the pH values of the samples would be
reduced to different levels below the threshold, as indicated by
the buffer curves obtained on the 1:5 suspension. The pH values
actually obtained—namely, 7.47 with the 0.5IV solution and 7.50
with the normal acid solution—are identical within limits of
precision of the pH determination. Obviously the soil at 70 per
cent of its water-holding capacity is buffered considerably more
than in its 1:5 suspension. This experiment is, nevertheless, of
interest in that it shows how reproducible the extent of microbiological nitrogen transformations actually is when the pH value
is held constant.
The resulting curves for these two series of incubations (D and
E), shown in Figure 4, are found to be so nearly identical that
they are practically superimposable. The numerical data for
this experiment, given in Table 3, indicate that the accumulation
of nitrite begins considerably sooner than in the B and C series.
Similarly, the period of time during which nitrites were produced
in appreciable amounts is considerably shorter in the acidified
than in the more alkaline samples. Nitrates were first observed at
a much earlier date in the acid-treated samples; hence by the
twenty-fifth day a considerably greater amount of nitrate had
accumulated in the D and E series than in the B and C series,
where the formation of nitrate was retarded by the high initial
The dependence of nitrate formation upon the attainment of the
threshold pH value of 7,7 is strikingly shown by the curves in
Figure 4. The control series A, in which the amounts of nitrogen
originally present in the samples were rather minute, could not
be expected to show a correlation between pH value and nitrate
formation. It is quite significant, however, that nitrates were
entirely absent from the samples of the B and C series until the
pH had dropped to the threshold value, after which time nitrates
began to accumulate. The importance of this threshold value is
further illustrated by the data in Table 3, which show that nitrate
formation began relatively promptly—i.e., on the sixth day after
the beginning of the incubation—in those samples whose initial pH
values were below the threshold. Therefore the conclusion may
be drawn that, if the initial pH value lies above the threshold,
nitratification begins promptly when the threshold value is
reached, and hence represents a response of the microorganisms to
an environment favorable to their activity. In samples whose
initial pH value lies below the threshold, nitratification appears
to depend simply upon the beginning of an appreciable microbiological activity following the normal lag phase. Since, in the
D and E series the initial pH values were close to the threshold
and during most of the incubation period did not drop to any
considerable extent below it, the foregoing deduction is consistent
with the normal growth or activity curve, the character of which
is indicated graphically in Figure 1.
The influence of the various treatments upon the initial and
final pH values of the soil samples and their relation to nitrite
and nitrate formation are shown in the form of a bar graph (Fig.
5). In this figure are also given the pH values at the time nitrites
were first observed, when the amount of nitrite formed reached
a maximum, and when nitrates made their first appearance. At
the top of each bar is given the day of incubation on which the
observation was made.
pH 25th. DAY
NH4OH - 0.5 N
Figure 5. — Threshold pH range in the nitrification of ammonia in Gila
sandy loam. (Numbers above bars indicate time in days of incubation after
which measurement was made.)
It will be noted in each instance that the pH value was equal to,
or less than, the threshold value before nitrates began to appear.
The time interval between the beginning of the incubation and the
first appearance of nitrates appears to vary directly with the
initial pH value of the sample. Thus in the most alkaline series
(B), 25 days had elapsed before nitrates appeared, and in the less
alkaline ammonia-treated series (C), 19 days elapsed prior to the
formation of nitrates. In the two acidified series (D and E), the
corresponding time interval was only 8 days. It is obvious, therefore, that the rate of nitrification of ammonia in desert soils
among other things is a function of the pH value, and, in general,
the higher it is above the threshold value of 7.7 =t 0.1 the longer
will be the time which elapses before nitrates begin to make
their appearance.
Although it was not the primary object of this investigation to
study the influence of pH upon the formation of nitrites in the
micr'obial oxidation of ammonia, sufficient evidence has been
accumulated to justify the formulation of a provisional rule as
follows: The more alkaline the soil, the longer will be the time
which elapses during the incubation before nitrites appear in
amounts detectable by analysis. Figure 5 also indicates that a
large decrease in alkalinity occurred in the B and C series prior
to the formation of nitrites. This decrease amounted to 1.4 pH
units in the B and 0.9 pH unit in the C series. The fact that complete equilibrium may not have been attained within the soil
samples can account for only a small portion of this decrease.
Hence it appears obvious that the ammonia must have been
oxidized to some intermediate product (or products) before passing over into the nitrite form, in order to account for the observed
drop in pH value. Further reference will be made to this phenomenon in the discussion of results which is to follow.
Another point of interest shown in Figure 5 is that the build-up
of nitrites appears to be conditioned by the length of time which
elapses, prior to the attainment of the threshold value, before the
oxidation of nitrite to nitrate begins. For example, in the B series,
nitrites were first observed on the seventeenth day of incubation;
nitrates on the twenty-fifth day. Thus there was a period of 8
days during which the accumulation of nitrites in the soil samples
could take place. By way of contrast, the corresponding time interval in the C series was 5 days, and in the D-E series, only 2 days.
While nitrites tend to accumulate in greatest amounts under
alkaline conditions, it will be seen from Table 3 that even in the
acid-treated samples, fairly large quantities of nitrite may accumulate, of the order of 70 to 90 p.p.m. Considering the process
quantitatively one may regard the total amount of nitrite
formed in the oxidation of ammonia as determined by two factors:
the rate of its build-up and the time during which such an accumulation can occur. This time interval is, as shown above, the
length of time required to bring the pH value of the soil down
to the threshold value, at which point nitrites would begin to
disappear by their oxidation to nitrate. It appears that high
alkalinity reduces the activity of the organisms, which is equivalent to reducing the rate of nitritification. It is quite possible that
ammonia is oxidized to nitrite more rapidly at the lower pH
values than under the more alkaline conditions and at a more
rapid rate than it can be oxidized to nitrates by the nitratifiers.
In the preceding experiments the incubations were set up and
begun under a particular set of conditions, and the pH value of
the soil in each case changed during the period of incubation. It
seemed of interest therefore to carry out an incubation experiment
in which the pH value of the soil, initially at a level considerably
above the threshold due to the addition'of ammonia, was maintained at a relatively high level by the periodic addition of base
in the form of calcium hydroxide solution. As in previous series,
100-gram samples of the cultivated Gila sandy loam were treated
with ammonia solution, added in 5-ml. portions to successive
quarter portions of soil, so that the total nitrogen added amounted
to 30 mg. The samples after weighing were incubated without
the close-fitting covers, so that evaporation would be hastened
and restoration of moisture lost could be made with saturated
calcium hydroxide solution. Determinations of pH, nitrate, and
nitrite were made on duplicate samples as before.
The data are shown graphically in Figure 6. The first samples
were analyzed on the third day of incubation, at which time the
pH value was 8.72. During the following 18 days, the pH value
gradually dropped in spite of the added base, but the nitrate accumulations were relatively slight, amounting to only 12 to 15
p.p.m. On the twenty-second day, however, notwithstanding the
pH value was still 7.90, the nitrate content had increased to nearly
35 p.p.m. This amount is small, to be sure, in comparison with
the amount formed in previous experiments, but nevertheless
significant, in view of the fact that the pH value was still slightly
above the threshold value. This result was interpreted to mean
that the pH-nitratification curve did not jail immediately to zero
once the threshold pH value had been passed, "but that nitratification continued to proceed slowly and nitrates to accumulate in
measurable amounts after the lapse of a sufficient length of time.
It might be taken to indicate that the pH value within a part
of the soil mass was less than the threshold value, which would
enable part of the nitrifying bacteria, at least, to function
Since the nitrites appeared to have reached a stationary value,
at least did not appear to increase or decrease, after the seventeenth day of incubation, and nitrates appeared to be forming
slowly within at least a portion of the soil mass, it was decided to
withhold additional base until the threshold value had been
reached, and then to observe whether or not such treatment had
favored the formation of nitrates. Figure 6 shows the result: On
the twenty-sixth day the pH value had decreased to a point within the threshold range—namely, 7.68—and the nitrate content had
risen to 75 p.p.m. On the thirtieth day, when the pH value was still
within the threshold range, the amount of nitrate formed had more
than doubled in amount, totaling 180 p.p.m. Evidently, after the
threshold pH range has once been attained, the nitrate content of
the soil increases normally as a result of ameliorated pH conditions, and this increase in nitrate content appears to be independent of the time interval which elapses prior to the attainment
of the threshold value.
Figure 6.—Rate of nitrification of ammonia in Gila sandy loam as influenced
by controlled alkalinity.
If high alkalinity is the factor which prevents nitratification, it
might be expected that the saturated calcium hydroxide solution
added to the surface of the soil samples to restore moisture lost
by evaporation would retard or prevent the activity of the nitrifying organisms with which it came in contact. It is probable,
however, that this treatment did not affect all of the organisms
throughout the sample equally, since there was no means available to bring about a uniform distribution of the base. The appearance of nitrate under these conditions can therefore be accounted for only on the assumption that the organisms were not
all affected in the same manner or to the same extent by the
calcium hydroxide solution added.
The experiment here described was, however, continued by
restoring the pH after the thirtieth incubation day to a value above
the threshold to determine whether, in line with the foregoing
reasoning, the rate of accumulation of nitrate should decrease.
Figure 6 shows, in fact, that the slope of the nitrate curve dropped
off quite promptly indicating that nitrates form much less rapidly
when the pH value of the soil is raised above the threshold value.
This experiment was terminated after 52 days of incubation, at
which time the nitrate content of the soil was found to have
reached a value of 225 p.p.m. This result is rather striking in
view of the fact that the same soil fertilized with ammonia in the
absence of calcium hydroxide had by the forty-ninth day of in-
cubation (as shown in Table 2) attained a nitrate concentration
of approximately 300 p.p.m. The nitrite curve in the same figure
illustrates the general result, observed in all of the experiments,
that the nitrite content of the soil falls rapidly to a mere trace as
soon as the threshold pH value is reached and nitrates have begun
to form at a fairly rapid rate.
In the foregoing studies several new and highly significant facts
have been brought to light: first, ammonia exhibits a nitrification
rate similar in magnitude to that of ammonium sulfate and urea
under similar experimental conditions; second, there exists a
threshold pH value for the nitratification of these compounds—
namely, 7.7±0.1—above which nitrites are not oxidized to nitrates
to any appreciable extent; third, nitrites may form in considerable
amount even in a well-aerated desert soil maintained at the
optimum moisture content.
The fact of outstanding interest is the existence of a threshold
pH value for nitratification, particularly in view of the work of
Gerretsen (24), Gowda (25), Meyerhof (37, 38, 39), Olsen (44),
and others, who reported pH values or ranges considerably higher
than 7.7 as the optimum for the nitrifying bacteria. Meyerhof
placed his optimum between 8.3 and 9.3, basing his deductions on
results obtained with pure culture media. Under such conditions
the mineral nutrients required by the organisms are usually
present in an easily available form, so that the bacteria should be
able to work most efficiently. In the soil, however, such may not
be the case. If the needed mineral nutrients are present in the
soil in a less available form due to high alkalinity, the bacteria
cannot function efficiently, if at all, in the nitrifying process until
the pH value has been reduced to a point where the required
nutrients are more readily available. Thus the threshold value
may prove to be a factor which affects both the biological behavior of the nitrifying bacteria and their nutrition as well.
The threshold value of 7.7±0.1 here found is in striking agreement with the value of 7.6 established by Breazeale and McGeorge
(11), beyond which plants are unable to absorb phosphate and/or
nitrate from the soil solution. Basing his deductions upon thermodynamic principles, Buehrer (13) showed that since plants are
unable to utilize phosphate from solution at pH values above 7.6,
the form of phosphate ion used by them must be the H2PO4~ ion.
Similarly Greene (26) found that Azotobacter grew best and
fixed nitrogen most abundantly in soils at reactions where phosphorus is present chiefly as the H2PCV ion, and concluded that
certain bacteria behave like plants with respect to their absorption of phosphate. It is therefore very probable that in alkaline
calcareous soils the nitratifiers are not able to perform their normal functions without phosphorus, for at reactions above the
threshold value the amount of phosphorus present in the form of
H2PO4~ ion is relatively small. In solution cultures, on the other
hand, it is quite possible for nitratification to occur at pH values
considerably higher than 7.7, in view of the fact that even at
higher alkalinity the ionic relations of the phosphate equilibrium
(see Table 3 and Ref. 13) permit the existence of an amount of
H2P04~ sufficient for bacterial metabolism. It is evident that the
percentage of the gross phosphate concentration represented by
the H2P04~ ion decreases as the pH rises. In solution cultures
where the total concentration of soluble phosphate is usually high,
the concentration of H2PO4" ions notwithstanding the high pH
value may still be sufficiently high to enable the nitrifying organisms to function. In the soil, however, where both the total
concentration of soluble phosphate and the percentage of H2PO4"
present at higher pH values are small in magnitude, the actual
concentration of H2PO4~ ion may be too small to satisfy the bacterial requirements.
It is a fairly well-established fact that the nitrifying organisms
can adapt themselves to free alkalinity if exposed to such conditions over sufficiently long periods of time. Olsen (44) working
with strongly acid humus soil adjusted the pH with lime and
found nitrification to occur over an unusually wide range of pH
values. It is quite probable that under these conditions he did not
succeed in neutralizing the acidity within the soil aggregates even
though the pH value, as determined on the entire soil mass, appeared to have been on the alkaline side of neutrality.
Among the nitrification studies on ammonia, ammonium sulfate, and urea, reported in the foregoing sections of this bulletin,
three observations are of particular interest: first, the failure of
the ammonia to nitrify in Superstition sand, as shown in Table 2;
second, the apparent irregularity in the nitratification occurring
in the virgin Gila sandy loam samples that had been fertilized
with ammonia, as shown in Table 2; and third, the slower rate
of nitratification of all three of the fertilizers in Palos Verdes
sandy loam, as illustrated by the data of Table 3.
The application of ammonia to Superstition sand, which exhibited a low buffer capacity toward base, being highly calcareous,
raised the hydroxyl-ion concentration to such an inordinately high
value—550 times the alkalinity existing at neutrality, for example—that nitrate formation did not take place. In the Palos
Verdes soil, however, the buffer conditions are reversed, the soil
having a low specific buffer capacity toward acid. This buffer
capacity is characteristic of acid soils of the type with which
Waksman (54) worked. He recommended the addition of a base
to neutralize the acid which is continuously formed during the nitrification. The Palos Verdes soil had an initial pH value of 6.5
and exhibited a characteristically slower rate of nitrate formation
than is shown by soils of a more alkaline reaction. During the
nitrification of ammonium sulfate in this soil, the rate of nitrate
formation decreased simultaneously with the decrease in pH value.
This retardation of nitratification rate at the lower pH values is
consistent with the findings of Waksman (54), Naftel (42),
Gaarder and Hagem (23), and Humfeld and Erdman (28). In the
present investigation, the series of samples treated with ammonium sulfate attained a minimal pH value of 4.65, which is not low
enough to verify the pH range of 3.9 to 4.4 adopted by the foregoing investigators as the extreme lower limit of pH for nitrate
The challenging aspect of the present series of experiments was
to establish definitely the actual existence of a threshold pH value
for nitratification, and particularly to determine whether the decrease in pH from the high initial values was accompanied by, or
resulted from, the formation of nitrites. The data for Series B
and C, as well as their graphical representation in Figure 5,
showed that a large drop in pH occurred prior to the beginning of
nitrite formation. Several factors may have contributed to this
effect: one is the possibility of a lag in the establishment of
equilibrium in the soil samples following application of the solutions; some time is manifestly required to establish moisture- and
base exchange-equilibrium in the soil. The effect manifests itself
as an almost immediate initial drop in the pH curve. It is a wellknown fact that the pH value of any soil tends to fall for some
time after the addition of water while the various equilibria are
being established, but the lowering of the pH value from this
cause is small in comparison with the large decreases in pH which
were observed prior to the beginning of nitrite formation.
Another factor which may have a bearing upon this decrease in
pH prior to nitritification is the fact that the first observations of
nitrite, as shown in Series B and C of Table 3, were not obtained
immediately after the formation of nitrites had begun. The first
appearance of nitrites in the B series was observed on the seventeenth day, at which time about 30 p.p.m. of nitrite had accumulated, or nearly half of the maximum amount which finally
formed. Similarly about 22 per cent of the maximum value for
nitrites in the C series had made its appearance by the fourteenth
day, when nitrites were first detected. If this reasoning is correct,
the remaining 50 per cent of the total amount of nitrite accumulated in Series B and 78 per cent in Series C was evidently much
less effective in reducing the pH value than the amounts mentioned above. Each p.p.m. of nitrite formed would, if one assumes
the equation
2 NH8 + 3 O2 + 2 OH- = 2 NO>~ + 4 H2O
to represent the process, remove an equivalent amount of hydroxyl ion; hence there is no justification for the assumption that
there may be a change in ratio of nitrite formed to hydroxyl
removed. It is apparent that there must be some intermediate reaction which is responsible for the lowering in pH prior to the
formation of nitrites.
Various investigators have assumed the presence of one or more
intermediate compounds in the formation of nitrite from am-
monia. Beesley (7), Mumford (40), Kluyver and Donker (30),
and Corbet (16, 17, 18, 19) have produced direct or indirect evidence of such intermediate products in their oxidation studies.
Beesley found that as much as 44 per cent of the applied ammonia
disappeared prior to the formation of nitrite, and concluded that
the nitrogen must have passed through some intermediate stage,
"which must be regarded as more or less hydroxylated." Mumford
advanced the theory that the oxidation of ammonia involves the
successive hydroxylation of the hydrogen atoms with its attendant
removal of water, and reported the presence of hydroxylamine
and salts of hyponitrous acid.
Kluyver and Donker (30) similarly incline to the hydroxylamine-hyponitrous acid mechanism for the oxidation of ammonia,
assuming that these intermediate products result from alternate
hydration and dehydrogenation characteristic of microbiological
reactions. Corbet (17) definitely reports the presence of these
two compounds in systems involving ammonia oxidation. He
writes as follows:
It appears that the reaction proceeds through the formation of hydroxylamine and hyponitrous acid as intermediate compounds, but while the
first-named can never have more than an ephemeral existence under the
experimental conditions, hyponitrous acid present in the form of the
calcium salt may account for as much as 40% of the total nitrogen present.
The actual presence of hydroxylamine and hyponitrous acid in
such a process is difficult to prove analytically even under the
most favorable conditions. Rao et al. (47) attempted to demonstrate the formation of these compounds by using culture technique, but without success. They are of the opinion that Corbet's
analytical procedure is not specific for hyponitrites. The presence of hydroxylamine, on the other hand, has been suggested
(8, 9, 10, 31, 33) as an intermediate product not only in the oxidation of ammonia to nitrite but also in denitrification and nitrogen fixation—both of which are microbiological processes.
The pH changes associated with the nitrogen transformations
in the microbiological oxidation of ammonia, set forth in the preceding sections, can be reconciled in part with the hydroxylaminehyponitrous acid mechanism as proposed by Kluyver and Donker.
To be tenable, such a mechanism must involve not only the removal of hydroxyl ions, which is responsible for the drop in pH
value, but it must also be consistent with the free energy changes
accompanying the respective steps in the process. The free energy
change, by virtue of its magnitude and sign, offers a criterion of
whether or not the postulated intermediate reactions are likely to
take place, or can take place if conditions are favorable. In other
words, it is a measure of the tendency with which the reaction
tends to take place, but it also represents the maximum amount
of available energy liberated or absorbed when one mol of the
reacting constituent is converted reversibly to another form.
Baas-Becking and Parks (5) employed such a free energy approach in a theoretical study of the energy efficiency of various
types of autotrophic bacteria, including the nitrifiers. They calculated the free energy change for the overall transformation of
ammonium ion to nitrite, and of nitrite to nitrate, basing their
calculations on certain data of Meyerhof (37, 38, 39) for the
optimum concentrations of the ions involved in the respective
transformations. The data in question were obtained in experiments with nutrient solutions, and from the analytically determined ion activities it was possible to calculate the free energy
change attending the process under those conditions.
In applying such free energy methods to transformations occurring in the soil, however, it is difficult to convert the standard free
energy changes to the actual activities of the constituents as they
exist in the soil solution, because neither hydroxylamine nor
hyponitrous acid can be satisfactorily determined by existing
analytical methods. It was decided, therefore, as a first approximation, to calculate the free energy changes for the constituents
considered as being present in their standard states—namely, at
unit activity. The appropriate data for these nitrogen transformations have been assembled by Latimer (32) from the most reliable
thermodynamic studies available.
In making such free energy calculations it is obviously possible
to base them either (1) on the half reaction involving the nitrogen compounds or ions, which would take place at the negative
electrode of a reversible galvanic cell, or (2) on the entire oxidation reaction involved in each step. In the latter procedure one
must introduce the oxidizing agent which is generally written as
molecular oxygen but which probably involves the microorganism in some manner. The free energy data so calculated are
shown in Table 4, in which the values are given both for the
half reaction and the entire oxidation reaction. For comparison,
the data have been calculated for the respective steps under both
alkaline and acid conditions.
It will be noted that when the nitrification of ammonia occurs
under alkaline conditions, as in the soils used in this study, each
of the steps (as shown by the equation for the half reaction) uses
up hydroxyl ions. As a result, there should be a decrease in pH
value resulting from each of these steps. Furthermore, there is
a much greater consumption of hydroxyl ion—namely, seven
mols—when one mol of ammonium ion goes to nitrite, than when
the nitrite is transformed to nitrate, which requires only two
mols. The fact that the hydroxylamine and hyponitrous acid steps
alone involve a total of five mols of hydroxyl accounts in part for
the phenomenal drop in pH which occurs prior to both nitrite and
nitrate formation.
The free energy changes attending these steps give a hint as to
a justification for the proposed mechanism. It will be noted that
the first step in which the hydroxylamine is formed involves a
large positive free energy change, "indicating that the process is
not spontaneous. If the reaction occurs at all, energy must be
supplied from some source, possibly by the microorganisms.
Moreover, the high positive value of the free energy suggests that
the hydroxylamine must be unstable, which in turn may account
fof the difficulty of identifying it in the soil. When it changes to
hyponitrous acid, the process is evidently spontaneous, since it
involves a considerable decrease in free energy (33,680 calories
per mol for the half reaction and 51,870 calories per mol for the
entire reaction). The fact that energy must be put into the system
in Step 1 may also account in part for the lag period which is always observed in the initial stages of ammonia oxidation.
The free energy data for the transformation of nitrogen under
acid conditions are of interest in the respect that in every step a
high positive free energy value is involved. The reaction is quite
evidently not favored by acid conditions since hydrogen ion is
formed in each step, which would tend to reverse the process. In
the case of the nitrifying organisms it has in fact been found (6)
that microbiological oxidation processes cease when the pH value
drops to 5.5, below which the oxidation is believed to be primarily
chemical and catalyzed perhaps by hydrogen ion. Below a pH
value of 4.5, the oxidation processes represented in Table 4 cease
entirely. It is of interest therefore to note that the oxidation of
ammonia, which is of primary importance when used as a fertilizer, proceeds favorably on the alkaline side of neutrality,
where the free energy values for the reactions are dominantly
negative. The fact that it requires the catalytic influence of the
microorganisms to proceed emphasizes the repeatedly observed
fact that the oxidation will not proceed with maximum efficiency
if the environment of the soil is not favorable to the metabolism of
the bacteria and their multiplication. Therefore, although an
excessively alkaline condition in the soil would theoretically
favor the nitrogen transformations in question, it is evident that
unless man can otherwise ameliorate such an adverse condition,
it must be done by the bacteria themselves, by reducing the pH
value to the point where their own nutrition processes become
normal and nitratification can take place efficiently.
1. A comparative study has been made of the rates of nitrification of ammonia, ammonium sulfate, and urea in six typical
Arizona soils.
2. A threshold pH value of 7.7±0.1 has been found for the
nitratification of the ammonia type of fertilizers in desert soils
above which the complete oxidation of ammonia will not occur,
and to which the pH value of such soils must first be reduced before nitrification can proceed to completion.
3. In the microbiological oxidation of nitrogen applied in the
three above-mentioned forms, there is considerable nitrite formation, even in well-aerated soils under favorable conditions of
temperature and moisture, so long as the pH value of the soil is
considerably above the threshold value. It was greatest in those
cases where the soil had been rendered strongly alkaline by addition of calcium hydroxide.
4. A pronounced decrease in pH value occurs in the soil prior
to both nitrite and nitrate formation.
5. Nitrite accumulation appears to be inhibited by a high concentration of calcium ions and/or by high alkalinity.
6. In all instances in which a significant accumulation of nitrites was noted, the amount of nitrite decreased almost simultaneously with the formation of nitrates.
7. Ammonia is not toxic to the nitrifying organisms even at a
concentration as high as 300 p.p.m. The failure of the ammonia to
nitrify in some instances is attributed to the high alkalinity of the
8. The fact that practically all of the nitrogen added as ammonia can be analytically accounted for indicates that losses by
volatilization from the soil or by the spontaneous decomposition
of ammonium nitrite are negligible.
9. Under constant (uniform) experimental conditions it is
found that equal amounts of nitrate are formed in soil samples
treated with the different nitrogen fertilizers after a given period
of time.
10. In alkaline calcareous soils ammonia nitrifies as rapidly as
ammonium sulfate and urea, provided the soil is sufficiently well
buffered toward base to withstand the initial change in pH. By
virtue of the microbial oxidation, the pH value of a soil treated
with ammonia will be reduced to substantially the same limiting
value as when an equivalent amount of ammonium sulf ate or urea
has been added to supply the nitrogen.
1. Arnon, D. I. 1937.
Ammonium and nitrate nitrogen nutrition of barley at different seasons in relation to hydrogen-ion concentration, manganese, copper
and oxygen supply. Soil Sci. 44:91-113.
2. Arlington, L. B., and Shive, J. W. 1935.
Rates of absorption of ammonium and nitrate nitrogen from culture
solutions by ten day-old tomato seedlings at two pH levels. Soil Sci.
3. Association of Official Agricultural Chemists. 1935.
Official and Tentative Methods of Analysis. Fourth Edition. A.O.A.C.,
Washington, B.C.
4. Ayers, A. D., and Jenny, H. 1939.
Private communication to Shell Chemical Company, San Francisco,
5. Baas-Becking, L. G. M., and Parks, G. S. 1927.
Energy relations in the metabolism of autotrophic bacteria. Physiol.
Rev. 7:85-106.
6. Barritt, N. W. 1933.
The nitrification process in soils and biological filters. Ann. Appl.
Biol. 20:165-76.
7. BeesleyrR. M. 1914.
Experiments on the rate of nitrification. Trans. Chem. Soc.
8. Blom, J. 1928.
Determination of hydroxylamine. Biochem. Zeit. 194:385-91.
9. Blom, J. 1928.
Formation of hydroxylamine in the reduction of nitrates by microorganisms. A contribution to the problem of amino acid formation by
microorganisms. Biochem. Zeit. 194:392-409.
10. Blom, J. 1931.
An attempt to clarify the chemical processes involved in the assimilation of molecular nitrogen by microorganisms. Centralbl. Bakt.
11. Breazeale, J. F., and McGeorge, W. T. 1932.
Nutritional disorders in alkaline soils as caused by deficiency of
carbon dioxide. Ariz. Agr. Exp. Sta. Tech. Bull. 41.
12. Buchanan, R. E., and Fulmer, E. I. 1928.
Physiology and Biochemistry of Bacteria. Williams and Wilkins Co.,
Baltimore. Vol. I, p. 16.
13. Buehrer, T. F. 1932.
The physico-chemical relationships of soil phosphates. Ariz. Agr.
Exp. Sta. Tech. Bull. 42.
14. Clark, H. E., and Shive, J. W. 1934.
The influence of the pH of a culture solution on the rates of absorption of ammonium and nitrate nitrogen by the tomato plant. Soil
Sci. 37:203-25.
15. Clark, H. E., and Shive, J. W. 1934.
The influence of the pH of a culture solution on the assimilation of
ammonium and nitrate nitrogen by the tomato plant. Soil Sci.
16. Corbet, A. S. 1934.
The formation of hyponitrous acid as an intermediate compound in
the biological or photochemical oxidation of ammonia to nitrous acid.
I. Chemical reactions. Biochem. Jour. 28:1,575-82.
17. Corbet, A. S. 1935.
Biological Processes in Tropical Soils with Special Reference to
Malaysia. W. Heffer and Sons, Ltd., Cambridge.
18. Corbet, A. S. 1935.
The formation of hyponitrous acid as an intermediate compound in
the biological or photochemical oxidation of ammonia to nitrous
acid. II. Microbiological oxidation. Biochem, Jour. 29:1,086-96.
19. Corbet, A. S. 1936.
The biological and chemical oxidation of ammonia to nitric acid.
Proc. Third Int. Cong. Soil Sci. 1:133-34.
20. Davidson, O. W,, and Shive, J. W. 1934.
The influence of the hydrogen-ion concentration of the culture solution upon the absorption and assimilation of nitrate and ammonium
nitrogen by peach trees grown in sand cultures. Soil Sci. 37:357-85.
21. Fraps, G. S., and Sterges, A. J. 1933.
Causes of low nitrification capacity of certain soils. Soil Sci. 34:35363.
22. Fraps, G. S., and Sterges, A. J. 1937.
Basicity of some phosphates as related to nitrification. Jour. Amer.
Soc. Agron. 29:613-21.
23. Gaarder, T., and Hagem, O. 1922-23.
Nitrification in acid solutions. Bergens Mus. Aarbok.; Naturv. Raekke
No. 1, 26 pp. Original not seen; Chem. Abs. 19:2,508 (1925).
24. Gerretsen, F. C. 1925.
Over den invloed van de waterstof-ionen-concentratie op bacteriologische processen. VersZagen van Landbouwkundige Onderzoek.
der Rijkslandbouwproefstat. 30:1-44. (Summary in English.)
25. Gowda, R. N. 1924.
Oxidation of ammonia and nitrites by microorganisms under different conditions. Soil Sci. 17:57-64,
26. Greene, R. A. 1933.
Some factors limiting the applicability of biological methods for
determining the availability of plant food elements in calcareous
soils. Soil Sci. 36:261-66.
27. Halversen, W. V. 1928.
The value of nitrification tests on soils representing extreme contrast
in physical and chemical properties. Soil Sci. 26:221-31.
28. Humfeld, H., and Erdman. L. W. 1927.
The significance of the hydrogen-ion concentration in soil nitrification studies. Iowa Acad. Sci. Proc. 34:63-67.
29. Jones, C. D., and Skinner, C. E. 1926.
Absorption of nitrogen from culture solutions by plants. N. J. Agr.
Exp. Sta. Ann. Rept pp. 360-65.
30. Kluyver, A. J., and Donker, H. J. L. 1926.
Die Einheit in der Biochemie. Chem, d, Zelle u. Gewebe 13:134-90.
Original not seen; cited by Stephenson (51).
31. Kostychev, S., and Tsvetkova, E. 1920.
The utilization of nitrates by molds for the production of nitrogenous
compounds. Zeit. physioL Chem. 111:171-200.
32. Latimer, W. M. 1938.
The Oxidation States of the Elements and Their Potentials in
Aqueous Solutions. Prentice-Hall, Inc., New York.
33. Lindsay, G. A., and Rhines, C. M. 1932.
The production of hydroxylamine by the reduction of nitrates and
nitrites by various pure cultures of bacteria. Jour. Bact. 241:489-92.
34. McGeorge, W. T. 1923.
The assimilation of nitrogen by sugar cane. Nitrates vs. ammonia
salts. Planters' Rec. 27:347-52.
35. McGeorge, W. T., and Martin, W. P. 1940.
pH determination of alkali soils. Jour, Assoc. Off. Agr Chem
36. Meek, C. S., and Lipman, C. B. 1922.
The relation of the reaction and of salt content of the medium on
nitrifying bacteria. Jour. Gen. Physiol. 5:195-204.
37. Meyerhof, O. 1916.
Untersuchungen iiber den Atmungsvorgang nitrifizierender Bakterien. Pflug. Arch, ges. Physiol. 164:353. Original not seen; cited
by Stephenson (51).
38. Meyerhof, O. 1916.
Untersuchungen uber den Atmungsvorgang nitrifizierender Bakterien. II. Beeinflussung der Atmung des Nitratbildners durch
chemische Substanzen. Pflilg. Arch. ges. Physiol. 165:229. Original
not seen; cited by Stephenson (51).
39. Meyerhof, O. 1917.
Untersuchungen liber den Atmungsvorgang nitrifizierender Bakterien. IV. Die Atmung des Nitritbildners und ihre Beeinflussung
durch chemische Substanzen. Pflug. Arch. ges. Physiol. 166:240.
Original not seen; cited by Stephenson (51).
40. Mumford, E. M. 1914.
The mechanism of nitrification. Jour. Chem. Soc. 30:36.
41. Naftel, J. A. 1931.
The absorption of ammonia- and nitrate-nitrogen by various plants
at different stages of growth. Jour. Amer. Soc. Agron. 23:142-58.
42. Naftel, J. A. 1931.
The nitrification of ammonium sulfate as influenced by soil reaction
and degree of base saturation. Jour. Amer. Soc. Agron. 23:175-85.
43. Nelson, D. H. 1929.
Some effects of manganese sulfate and manganese chloride on nitrification. Jour. Amer. Soc. Agron. 21:547-60.
44. Olsen, C. 1928.
On the significance of hydrogen-ion concentration for the cycle of
nitrogen transformation in the soil. Compt. rend. Lab. Carlsl>erg
17: (8): 21. Abstracted in Nature (London) 123:44 (1929).
45. Pierre, W. H. 1927.
Buffer capacity of soils and its relation to the development of soil
acidity from the use of ammonium sulfate. Jour Amer. Soc. Agron.
46. Prianishnikov, D. N. 1929.
Ammonia in fertilizers and its relation to the life of plants. Trans.
Sci. Inst. Pert. (Moscow) 61:99-103. Abstracted in Chem. Abs.
23:5,264 (1929).
47. Rao, W. V. S,, Krishnamurti, P. V., and Rao, G. G. 1938.
Mechanism of the microbiological oxidation of ammonia. I. Formation of intermediate products. Jour. Indian Chem. Soc. 15:599-603.
48. Sessions, A. C,, and Shive, J. W. 1933.
The effect of culture solutions on growth and nitrogen fractions of
oat plants at different stages of their development. Soil Sci
49. Stahl, A. L., and Shive, J. W. 1933.
Studies on nitrogen absorption from culture solutions. I. Oats Soil
Sci. 35:375-99.
50. Stahl, A. L., and Shive, J. W. 1933.
Further studies on nitrogen absorption from culture solutions
Buckwheat. Soil Sci. 35:469-83.
51. Stephenson, Marjory. 1939.
Bacterial Metabolism. Longmans, Green and Co., New York pp
52. Stewart, G. R., Thomas, E. C., and Horner, J. 1925.
The comparative growth of pineapple plants with ammonia and
nitrate nitrogen. Soil Sci. 20:227-42.
53. Tandon, S. P., and Dhar, N. R. 1934.
Influence of temperature on bacterial nitrification in tropical countries. Soil Sci. 38:183-89.
54. Waksman, S. A. 1923.
Microbiological analysis of soils as an index to soil fertility. V.
Methods for the study of nitrification. Soil Sci. 15:241-60.
55. Waksman, S. A. 1932.
Principles of Soil Microbiology. 2nd Ed. Williams and Wilkins Co.,
Baltimore, Md.
56. Waynick, D. D. 1934.
Anhydrous ammonia as a fertilizer. Calif. Citrograph 19:295.
57. Willis, L. G., and Piland, J. R. 1931.
Ammonium-calcium balance: A concentrated fertilizer problem
Soil Sci. 31:5-17.
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