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IV. Soil Phosphorus in the Solid Phase

IV. Soil Phosphorus in the Solid Phase

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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 '

Extraction number

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

meaning. Here




+ bP)

or in linear form

P / v = l/bv,

+ P/v,

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)







Lime Potential

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

of hydroxylapatite.

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

2 pOH)

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



2 :1

2 :1





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

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IV. Soil Phosphorus in the Solid Phase

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