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IV. Soil Phosphorus in the Solid Phase
were linear, penetration of phosphorus into some amorphous region of
the crystals occurred.
Studies of this kind are valuable in that they give some indication
of what may be occurring in soil. It must be remembered, however, that
only the lower phosphorus concentrations used in such studies are relevant to intact soil. The more concentrated solutions are relevant to soil
which comes into immediate contact with the solutions from dissolving
The results from pure systems may not strictly apply to the much
more complex environment of soil, Thus, when Fried and Shapiro (1956)
made repeated extractions of soil with distilled water, they observed
two different patterns of phosphorus release (Fig. 4). For soils 2, 3, 5,
0 ' 3
i 2 '
FIG.4. Phosphorus-release curves of six soils as measured by successive 1-hour
water extractions. (From Fried and Shapiro, 1956.)
and 7 (pH range 5.6 to 6.5) the phosphorus concentration in successive
extracts gave a pattern which they interpreted as being due to a
desorption-type reaction. For three other, more acid, soils (e.g., 8 F ) the
pattern of release was thought to be due to the dissolution of sparingly
soluble phosphorus compounds.
However, both patterns could well be due to the dissolution of phosphorus compounds. Hydroxylapatite, for example, might be expected to
show a similar release pattern to that given by the high pH soils. Such
a pattern was shown by Deitz et al. (1964), who explained their observation by the formation of a surface complex on the hydroxylapatite
Calo(POa)c(OH)z 6 H20
+ 2 HP042-
4 Ca2HPOr(OH)2 2 Ca2+
It may be seen that this reaction is in fact an incongruent dissolution in
that the Ca:P ratio in the solution differs from that in the solid. It may
also be argued that as hydroxylapatite itself has colloidal properties
(Mattson et al., 1951), its dissolution is a surface chemical phenomenon
which will resemble a desorption mechanism. Conversely, the precipitation of hydroxylapatite will show a similarity to adsorption.
Measurement of the size of the phosphorus adsorption system can be
made by fitting adsorption data to a previously described adsorption
isotherm. The Freundlich isotherm has often been used for this purpose
and has the form
or in linear form
logX = logu b logc
where X is the amount of phosphorus adsorbed per unit weight of soil,
c, is the concentration of phosphorus in solution and a and b are constants that vary between soils. This isotherm is purely empirical, and the
constants have no physical meaning.
In contrast, the Langmuir isotherm has constants which, at least
when applied to the adsorption of gases on solids, have a quantitative
or in linear form
P / v = l/bv,
where P is pressure, o the amount of gas adsorbed per unit weight of
solid, om the maximum amount of gas which can be adsorbed as a
monolayer, and b a constant related to the energy of bonding. The principal postulates of the Langmuir isotherm (Adamson, 1960) follow: (1)
the energy of adsorption is constant (which implies uniform sites and no
interaction between adsorbate molecules); ( 2 ) the adsorption is on
localized sites (which implies no translational motion of adsorbed molecules in the plane of the surface); ( 3 ) the maximum adsorption possible
corresponds to a complete monomolecular layer.
It seems unlikely that all these postulates will hold for phosphorus in
soil. For example, the energy of phosphorus adsorption is likely to be
constant only within a narrow concentration range. Also it seems unlikely
that adsorption will be restricted to a monolayer, particularly at higher
concentrations where some sort of lattice structure will begin to form.
It seems unlikely therefore that the Langmuir isotherm will be strictly
applicable to phosphorus adsorption by soil.
Olsen and Watanabe (1957), however, did find that adsorption followed the Langmuir isotherm. They used the phosphorus equilibrium
concentration ( c ) instead of the partial pressure ( P ) and obtained a
linear relationship between c/v and c for dilute solutions. Thus they were
able to calculate the two constants b, related to the bonding energy of
phosphorus to the solid, and vm, the maximum adsorption capacity. However, they stressed that the fit of their experimental data to the Langmuir
isotherm does not provide any information concerning the mechanism by
which phosphorus is retained.
In more concentrated solutions, giving equilibrium concentrations
greater than say l e 3 M , Olsen and Watanabe found that the Langmuir
plots were no longer linear. Furthermore, it is doubtful that their
observations even at lower concentrations are generally applicable. For
example, in this laboratory, phosphorus adsorption isotherms were determined for 120 soils, and the relationship between c / v and c was found to
be curvilinear in the maioritv of the soils even when thevJ were eauiliI
brated with very dilute solutions giving equilibrium values less than 6 x
10-4 M .
Despite these limitations the Langmuir isotherm can often be used to
give a relative measure of the energy by which phosphorus is bonded
to the solids and a relative adsorption maximum. Based on this adsorption maximum and the amount of readily isotopically exchangeable phosphorus already present in the soil, an arbitrary calculation of the degree
of phosphorus saturation can be made. Such a measure has been shown
to be related to plant uptake of soil phosphorus (Gunary and Sutton,
The adsorption of both orthophosphate and pyrophosphate by soils
has been measured in this laboratory, using a modification of the procedure of Olsen and Watanabe (1957). The adsorption maxima were
generally found to be higher for pyrophosphate than for orthophosphate
(for example Fig. 5 ) , even when allowance was made for the fact that
one pyrophosphate ion contains two phosphorus atoms (Sutton and
Larsen, 1964; Gunary, 1966).
These data imply that soil particles have adsorption sites which hold
pyrophosphate but not orthophosphate. This implication is unacceptable
when it is considered that pyrophosphate ions are larger than orthophosphate ions. Also, pyrophosphate ions are of higher valency and electrical
charge, which makes it difficult to explain the smaller bonding energies
of pyrophosphate compared with orthophosphate ( Sutton and Larsen,
1964) in terms of a simple anion adsorption mechanism, Just as with
cation adsorption, one should expect polyvalent anions to be retained
with greater energy than monovalent ones.
It is reasonable, therefore, to conclude that the phosphorus ions are
not taken up by the soil particles by a simple adsorption mechanism. By
assuming that adsorption (or just sorption) is a combination of precipitation and adsorption, the apparent discrepancies may be reconciled. For
orthophosphate, the apparent adsorption pattern may be due to the
formation of small microcrystals of hydroxylapatite, which adhere to
active soil surfaces. Since pyrophosphate has a very marked effect in
conc. (moles x1@/ liter)
FIG. 5. Adsorption isotherms for orthophosphate and pyrophosphate in two
contrasting soils. (From Cunary, 1966.)
preventing the normal precipitation of hydroxylapatite ( Fleisch et al.,
1965), it is reasonable to assume that pyrophosphate addition leads to
the formation of smaller and less stable microcrystals. These smaller crystals may also have a greater tendency to adhere to the surface of soil
particles and therefore increase the apparent adsorption maxima.
Needless to say, this hypothesis needs closer examination, preferably
using pure minerals as adsorbants, but it seems to offer a chance of
resolving the controversy between the adsorption and precipitation
Geochemically, phosphorus, unlike many other minor constituents of
the Earth's crust, is always mineral-forming and does not just substitute
in other lattice structures (Landergren, 1962). However, in soil, most of
the inorganic phosphorus occurs in the clay fraction from which it cannot
be separated by physical methods. Consequently, direct evidence of the
nature of the inorganic phosphorus cannot be obtained by known petrographic methods. Only when phosphorus has been separated from the
soil (or is formed in layers or pockets in the soil by natural processes)
can a sufficient concentration of phosphorus minerals be obtained for
direct petrographic examination. So far, only the phosphorus minerals
apatite, vivianite (Fe,( PO,),.8 H,O) and wavellite (A&( OH),( PO,),.
5 H,O) have been qualitatively determined in soil by such methods
( Black, 1957).
A semiquantitative method for the direct determination of soil apatite
has recently been developed by Shipp and Matelski (19m). They isolated the heavy minerals from soil samples and treated them with a few
drops of concentrated sulfuric acid. This caused needlelike calcium
sulfate crystals to develop on the surface of the apatite particles, which
were thereby recognized and counted. They observed that the number
of apatite particles increased with soil depth and reflected the distribution of acid-soluble phosphorus in the soil. Although this method provides
a direct way of detecting apatite minerals, it does not distinguish between
the various forms of apatite, e.g., fluoroapatite and hydroxylapatite. Since
no effervescence occurred it can be assumed that the apatite was not
formed on the surface of carbonate minerals, and that it did not contain
appreciable amounts of carbonate.
Attempts have been made to classify inorganic soil phosphorus into
different compounds, according to their extractability in various reagents
(Chang and Jackson, 1957a). Such methods must be arbitrary since the
reagents will inevitably cause a redistribution of the phosphorus during
the extraction. Thus, compounds reported to be present may not have
been so in the original soil.
Of the available evidence concerning the nature of phosphorus in
intact soil, the most relevant is that obtained from the application of the
solubility product principle. On this basis, the occurrence of phosphorus
compounds of calcium, aluminum, and iron have been suggested,
1. Calcium Compounds
Bassett (1917) stated than in pure systems hydroxylapatite is the
stable phase over a wide acidity range, and concluded that this is probably the only calcium phosphate that can permanently exist under normal soil conditions.
Nevertheless, determinations of the solubility of inorganic soil phosphorus in dilute salt solutions have led in one instance to the statement
that calcium phosphate, and by implication hydroxylapatite, was not
present in soil of moderate acidity (Clark and Peech, 1955), and in
another, that octocalcium phosphate ( Ca,H ( PO, ) 3 H,O ) was present
in soil which had been limed and fertilized with superphosphate ( Aslyng,
Octocalcium phosphate is not found as a mineral in nature, and it is
very difficult to make in the laboratory, where its formation has been
shown to be very sensitive to the presence of impurities ( Bjerrum, 1958).
The formation of octocalcium phosphate in soil therefore seems very
improbable. If it did occur, one would also have expected it to have been
observed, as it forms fairly large crystals which are easily detectable
even by a crude light microscope (Bjerrum, 1958).
Many assumptions about calcium phosphate minerals in soil are based
on insufficient understanding of the complicated solubility product of
hydroxylapatite, which varies by a factor of 10" (Van Wazer, 1958).
The influence of carbonate on the solubility of hydroxylapatite was discussed in Section 111, B, 1, where it was shown that the solubility of soil
phosphorus in 0.01 M CaCla over a range of pH levels paralleled that of
pure hydroxylapatite in the presence of carbonate. It will suffice here
to refer to a solubility diagram (Fig. 6 ) from which it may be seen that
there is no tendency for the points to be more numerous in the region
corresponding to octocalcium phosphate. The even distribution down to
a lime potential of about 3.5 and the absence of a higher density of points
in the region corresponding to hydroxylapatite, is in agreement with the
above effect of pH in the presence of carbonate.
Thus there is no reason to assume that any calcium phosphate other
than hydroxylapatite is permanently present in slightly acid, neutral, and
alkaline soils, and a knowledge of the properties of hydroxylapatite is
therefore essential for a better understanding of phosphorus in many
In a discussion of phosphorus in alkaline and calcareous soils, Olsen
(1953) gave a comprehensive review of the known properties of calcium
phosphates. In the intervening years the following main aspects of
hydroxylapatite have been studied: ( a ) formation from aqueous solutions, ( b ) crystal structure and the accommodation of impurities in the
lattice, ( c ) surface properties.
a. Hydroxylapatite formation. By means of X-ray diffraction analysis, Denk and Christensen (1962) studied the solid phase formed when
KOH was added to a CaC12 solution which contained excess KH,PO,.
They found the precipitated material to be crystalline dicalcium phosphate which gradually changed to hydroxylapatite when the pH of the
solution was between 6 and 14. The rate at which hydroxylapatite was
formed increased with pH. It would thus appear from their results that
dicalcium phosphate can act as a precursor for hydroxylapatite. Eanes
et al. (1965) also studied the intermediate stages in the precipitation of
hydroxylapatite, but found that the initial solid phase was noncrystalline
and the Ca :PO, ratio was near 1.5. Crystalline hydroxylapatite began to
form after about 5 hours, and this process was complete about 2 hours
later. After that time slow changes occurred in the crystalline structure,
so that a Ca:PO, ratio of 1.67 was eventually attained. The rate of transformation from noncrystalline material to crystalline hydroxylapatite
strongly suggested that the conversion mechanism was autocatalytic.
Taylor and Gurney (1965)
Clark and Peech (1955)
Clark and Peech (1960)
FIG. 6. Phosphate and lime potentials of soils in relation to those of pure
calcium phosphates. The lines for hydroxylapatite ( HA ), octocalcium phosphate
(OCP), and dicalcium phosphate (DCP) are based on the solubility data of Bjerrum
The authors did not name the noncrystalline material although the Ca:
PO, ratio would suggest a “tricalcium phosphate.”
Comparing the findings of Denk and Christensen with those of
Eanes et al., it should be borne in mind that the former used a solution
with an excess of phosphorus whereas the latter used a solution with an
excess of calcium, which is more realistic as far as soil is concerned.
b. Hydroxylapatite structure. Since the crystal structure of apatitic
calcium phosphates were independently described by Naray-Szabo
( 1930) and Mehmel (1931), several improved structural models have
been suggested, the latest being that of Kay et al. ( 1964). The illustration
of hydroxylapatite structure shown in Fig. 7 is that taken from Arnold
(1950) in which column calcium atoms are distinguished from those in
the lattice layers.
It appears that the calcium in hydroxylapatite cannot be substituted
by magnesium, and it has been suggested by Tovborg-Jensen and Rowles
( 1957) that magnesium ions inhibit the growth of hydroxylapatite nuclei
by being adsorbed onto the surface and blocking further lattice growth.
FIG. 7. Schematic xepresentation of apatite structure, oxygen atoms being
omitted. (From Arnold, 1950.)
In contrast, strontium ions can replace calcium in hydroxylapatite
and a strontium hydroxylapatite corresponding to calcium hydroxylapatite is known ( Niaki, 1961 ) .
Carbonate is the most important impurity in hydroxylapatite. It is
always found in the hydroxylapatite in sediments and even laboratory
preparations nearly always contain some, the quantity depending on the
amount of care taken in excluding COz during the preparation and aging
of the hydroxylapatite. For many years the form in which the carbonate
was present has been a matter for conjecture. Opinions differed as to
whether carbonate was present as calcium carbonate (calcite or aragonite
either adsorbed or as a separate phase), or alternatively, as a substitute for hydroxyl, giving a unit cell with the chemical composition
Calo( PO,) &03.
Recent studies of carbonate apatites by means of X-ray diffraction
analysis have conclusively shown that CO, is replacing PO, (ZapantaLeGeros, 1965; Smith and Lehr, 1966; Trueman, 1966). Smith and Lehr,
who investigated sedimentary apatites, suggested that one PO ion was
replaced by one CO, and one F ion. Thus a graph of PO, deficit against
CO, present should give a straight line with a slope of -1. They found,
however, a slope of -0.66 mole PO,/mole CO,. This slope agrees remarkably well with that of -0.67 found by Zapanta-LeGeros and reasonably well with that of -0.63 found in this laboratory. It should be noted
that these latter preparations did not include fluoride. It thus seems that
three CO, ions are replacing two PO, ions. The question of reconciling
this replacement with the structure of the apatite lattice falls outside the
framework of this article.
The effect of carbonate in the hydroxylapatite lattice is to make it
more chemically reactive, and it has been suggested (Larsen, 1966b)
that the electrical surface properties associated with colloidal hydroxylapatite (Mattson et al., 1951) are caused by this substitution.
Hydroxylapatite crystals occur as rods which in well crystallized
preparations have a length of about 200 A. and a width of 65 A. (Carlstrom, 1955). The particle size is thus very small and the specific surface
consequently very large, generally between 20 and 130 m.? per gram.
The surface properties are important .and often decisive for the chemistry
c. Surface properties of hydroxylapatite. Arnold ( 1950) suggested
from consideration of the imperfect unit cells on the surface of hydroxylapatite crystals, that a surface calcium phosphate existed with the chemical composition Ca2P0,0H. In a series of studies of the solubility product in hydroxylapatite/water systems, Rootare et al. ( 1962) suggested
that the hydrated form, Ca,HPO,( OH),, occurs as a surface complex on
hydroxylapatite crystals. They envisaged the following equilibrium:
Calo(P04)G(OH)2 6 HzO
+ 2 HP04Z-
4 C ~ Z H P O ~ ( O H 2) ~Ca2+
This reaction would give a Ca :PO, ratio of 1: 1in the solution, indicating
an incongruent dissolution.
When the surface compIex itself dissoIves
+ HP042-+ 2 OH-
C a ~ H p 0 4 ( 0 H2
) ~2 Ca2+
the cumulative ratio of Ca:PO, in the solution should approach 5 : s
(1.67), the ratio in hydroxylapatite. Rootare et al. demonstrated this in
their experiments. They also calculated the “solubility product” of the
surface complex and found the value pK‘ (2 pCa pHPO,
to be 27.3 at 25°C. and 25.1 at 40°C.
Francis (1965), in studies of the solubility behavior of dental enamel
and other calcium phosphates, used buffer solutions varying in pH from
5 to 6.3. She found that the Ca:PO, ratio in the solution, after allowing
for formation of calcium complexes, was greater than that in hydroxylapatite, indicating an incongruent dissolution in the pH range 4 to 6:
+ 6 HP04,+ 2 H20 WzO) 4 Ca2+ + 6 Ca HPO4.2 HzO
Ca~,(PO4),(0H), 8 H+ -+ 10 Ca2+
She concluded from solubility measurements that CaHPO, 2 H 2 0 was
found on the surface of synthetic hydroxylapatite, enamel, and bone.
In studies of human bone, MacGregor and Brown (1965) used the
relationship between lime potential and phosphoric acid potential to
characterize the chemical composition of the surface calcium phosphate
on hydroxylapatite. They found this to correspond to octocalcium phosphate in child bone and to hydroxylapatite in adult bone and suggested
that octocalcium phosphate acted as a precursor for formation of hydroxylapatite in bone.
Bjerrum ( 1958) considered the apparent adsorption of dicalcium
phosphate on the surface of hydroxylapatite to be a substitution of calcium by hydrogen ions. He suggested that a unit cell with the chemical
composition 6 CaHPO, 6 H 2 0 may be formed on the surface when
enough substitution had taken place, and he stressed the dependence of
the reaction on pH. Bjerrum also discussed the formation of a surface
complex of the composition Ca( H,PO,),; he regarded it as improbable
but not impossible.
To sum up, the calcium phosphates listed in the tabulation have been
suggested as occurring on the surface of hydroxylapatite.
2 4 . H20
Ca HP04. 2 HzO
(Ca HP04), . 6 HzO
The apparent discrepancies between the various results obtained may
be reconciled by assuming that all these surface complexes are possible
and that the boundary between them is d i h s e . This would lead to an
acceptance of the amphoteric property of hydroxylapatite suggested by
Mattson et al. (1951).
Further studies will no doubt provide more detailed information on
the surface properties of hydroxylapatite and thus, by implication, about
the peculiar solubility properties of this phosphate. There can be little
doubt that hydroxylapatite solubility is determined by the surface, not
by the deeper, layers of the crystal.
The understanding of phosphorus in a great many soils depends on
our precise knowledge of hydroxylapatite in pure systems, and although
progress in this field is being made all the time, there is still much to be
learned. Our present knowledge suggests that the formation of hydroxylapatite occurs through chemical changes in the solid phase, after an
initial precipitation of some undefined calcium phosphate has taken
The many possible surface complexes suggested for hydroxylapatite
underline what was already deduced from solubility data, that the
application of the solubility product principle to hydroxylapatite ( and
by implication to soil in which this phosphate may occur) is complicated.
A solubility product corresponding to a discrete calcium phosphate does
not imply that this phosphate is present as a mineral in its own right,
for it may merely occur as a surface complex on hydroxylapatite or even
on nonphosphatic minerals.
2. Aluminum Compounds
The occurrence of the aluminum phosphate mineral variscite with
the chemical composition AlPO,.2 HzO has been postulated in soil of
slight acidity, and the postulate has been supported by calculation of
the appropriate solubility product ( Chang and Jackson, 195713; Lindsay
and Moreno, 1960). However in pure systems, only where the pH of the
equilibrium solution is less than 3.1 does the solubility product of
variscite control the phosphorus concentration in solution (Bache, 1963).
At higher pH values, variscite dissolves incongruently, whereby a more
basic solid phase of aluminum hydroxyphosphate is formed (Taylor and
Gurney, 1962a,b, 1964). This material, by forming a surface complex on
variscite probably controls the phosphorus concentration in solution in
acid soils. Only when the surface of the variscite is very large, that is
when the crystal size is very small, is it likely that the complex constitutes
the bulk of the solid phase of aluminum phosphate which may be
Raupach (1963) has also considered the aluminum contents of displaced soil solutions and soil extracts in relation to pH and phosphorus
concentration. He found the results to be compatible with the existence
of ions such as A10H2+ and AI(OH),+ in solution, as found in pure
aluminum hydroxide/water systems, but he could find no evidence for
the existence of variscite in the soils studied. He concluded that the