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III. Phosphorus in Soil Solution

III. Phosphorus in Soil Solution

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SOIL PHOSPHORUS



155



liquid phases and also with the ratio between solid and solution. The

forms in which phosphorus exists in solution are governed by the reactions of protonation and complex formation. These forms are phosphoric acid, H,PO,, the corresponding ions, HLP04-,HP0,2-, and Po,”,

and the soluble complexes of these ions.



1. Protonation

The distribution of phosphorus between phosphoric acid and the

phosphorus ions is determined primarily by p H in the following equilibrium:

Peeid



*



PbsRe



+ Hf



From the law of mass action



where K is the dissociation constant of the acid and the brackets indicate

concentrations. This simple dissociation “constant” is influenced by the

temperature and concentration of the solution. The latter factor can be

accounted for by using activities, thus writing aH, etc., fbr active concentrations:

a--



&H



. &Pbnse



aPbase



Using the notation pH = -log aH the equation can be written:



Phosphoric acid is a tribasic acid, and it has consequently three dissociation constants from which the proportions of the free acid and the

phosphorus ions which exist at various pH levels can be calculated. The

values presented in Fig. 1 are based on the values of the dissociation



PH



FIG. 1. Distribution of ion species of phosphoric acid at different pH values at

infinite dilution (full line) and at ionic strength of 0.03 (broken line).



156



SICURD LARSEN



constants at 18°C. given by Bjerrum ( 1958):

pK’ = 2.120 - 0 499 4m -0.34~~

< f 1.04m

pK” = 7.227 - 1497 m

f 2.25m

pK”’ = 12.465 - 2.495 4%



where m = ionic strength.

It is obvious from the graph that the ionic species H,PO,- and HP04’are the most abundant in the pH range encountered in soil and that the

ionic strength encountered in nonsaline soil does not alter this distribution significantly. It should be remembered that increased ionic strength

normally depresses soil pH.

2. Complex Formation

Phosphorus forms soluble complexes with many metallic ions, and

part at least of the phosphorus in the soil solution may be complexed.

The soluble complexes that are likely to occur and their equilibrium

constants, expressed as stability constants, are given in Table 11. The

greater this constant, the more stable is the complex.

A normal soil solution will contain several kinds of metallic ions

which will form complexes with ions other than phosphate (e.g., hydroxyl, carbonate, sulfate, and organic ions). In addition, phosphorus

ions may take part in other homogeneous reactions. It would be a very

formidable task to calculate the proportion of the phosphorus that is

likely to be present in a complexed form in a given soil solution. Even

in a simplified soil solution dominated by one electrolyte it is still difficult

to make more than a rough approximation.

However, there has been some recent work in which the presence of

complex ions have been investigated in soil extracts. Weir and Soper

(1963) found that phosphorus was held in a soluble complex of ferric

humate formed when humic acid was extracted from an acid soil by a

solution of ferric chloride. They discovered that the amount of phosphorus in the complex decreased slightly as the p H of the solution was

raised from 5 to 9. Taylor and Gurney (1962a) concluded that an acid

suspension of colloidal aluminum phosphate contained significant quantities of a soluble Al-P complex. Fordham (1963) found no evidence for

the presence of complex phosphate ions in a 0.01 M CaC1, extract of a

soil of pH 6.3. However, in a calcareous soil Larsen (1965, 1966a) found

that the increase in the solubility of phosphorus, relative to the calcium

chloride concentration of the extractant, could be explained by assuming

the formation of the soluble complex CaHPO,. He found no such effect in

an acid soil of pH 5.5 where the HPO, ion concentration is negligible.

No information concerning soluble phosphorus complexes in the true



157



SOIL PHOSPHORUS



TABLE I1

Stability Constantsa of Metal-Phosphate Complexes

~



Cation



Reaction



Mgz+



+ HP042-+ NaHP04K+ + HP042- + KHPO4Mg2+ + HPOn2- C MgHPO4



Ca2+



Ca2+



Na+



Na+



K+



+ HPOa2-



CaHP04



Ionic

strength

(mole/l.)

0.2



Temp.



("C.)



0.2

Zero

0.16

Zero

0.2

0.2

0.006

Zero

0.15

Zero

0.2

0.1

0.1

0.1



0

25

0

25

25

25

38

38

25

25

22

25

37

25

25

18

18

18



0.4



25



0.4



25



0.2



0.665

Zero

Variable

a



~~~~



30

?

?



LoglOlC

0.08

0.60

0.08

0.49

1.88

2.50

1.62

2.87

1.50

1.70

2.20

2.70

1.86

1.08

2.58

ca. 3

ca. 2 . 3

ca. 2 . 3

3.49

13.45

8.23

j8.36

9.35

9.75

9.15



Compiled from Sillen and Martell (1964).



soil solution is available, but based on the meager experimental evidence

above and the known stability constants of phosphorus complexes (Table

11),it seems likely that a significant proportion of the phosphorus in the

soil solution may sometimes be present in this form. Further, it would be

expected that this proportion would be high both in acid soils (due to

the high stability of Fe-P and A1-P complexes) and in calcareous soils

(due to the high proportion of HPO, ions). Complex formation will be

lowest in slightly acid soils where H,PO, ions predominate and where

the concentrations of iron and aluminum are low.

Thus although our knowledge about soluble phosphorus complexes

in the soil solution is very limited, it seems an oversimplification to assume that even the bulk of the soluble phosphorus in soil is always present as the two ion species H2P0,- and HPO,*-.



158



SIGURD LARSEN



B. H E ~ O G E N E OEQUILIBRIA

US

The upper limit for the phosphorus concentration in solution is set by

the heterogeneous equilibria in which it takes part, The reactions involved are the dissolution and precipitation of sparingly soluble phosphorus salts, controlled by the solubility product principle and by the

adsorption of phosphorus on the surface of soil particles. Because the

phosphorus concentration in solution is governed by reactions with the

solid phase, the equilibrium level can be used to characterize the energy

state of phosphorus in the whole system. This has led to a considerable

amount of work on the chemical potential of phosphorus in soil.



1. Solubility Product Principle

The solubility product of a salt AaBb taking part in the following

equilibrium:

A,& G aA



+ bB



is defined as



When the salt is present in the solid phase, AaBb is constant and the

solubility product simplifies to

K q = [Ala[Blb



the value of which depends on the ionic strength of the solution. A true

constant can be obtained by using activities instead of concentrations,

and using the convention that p = --log,,

the solubility constant can

then be expressed:

PK., = a pA



+ b pB



The relevant sparingly soluble phosphorus salts are those of magnesium, calcium, aluminum, and iron. Solubility products for these have

been tabulated by Sillen and Martell ( 1964).

Wild (1954) examined a large number of early analytical data but

found no agreement between the concentration of phosphorus in the soil

solution and the concentration predicted from solubility products. More

recently, several other workers in this field (Chakravarti and Talibudeen,

1962; Hagin and Hadas, 1962; Bache, 1963) have also found that the

phosphorus concentration of the soil solution did not conform to solubility product principles.

In neutral and calcareous soils the absence of agreement between the

observed phosphorus concentrations and the solubility of calcium phos-



SOIL PHOSPHORUS



159



phates may be explained by the incomplete understanding of the solubility product of these compounds. Thus Bjerrum (1949) found two

solubility products for octocalcium phosphate and two for hydroxylapatite, one when the equilibrium was approached by precipitation and

the other when it was approached by dissolution.

A further complication is the strong influence that impurities seem

to have on the solubility of basic calcium phosphate. Thus Greenwald

( 1942) and Ericsson (1949) found that pure hydroxylapatite obeyed the

solubility principle but in the presence of small amounts of calcium

carbonate it did not. The equilibrium solutions were apparently supersaturated with respect to hydroxylapatite.

The deviation from pure hydroxylapatite behavior can be given by

the “saturation index” n suggested by Bjerrum ( Schmidt-Nielsen, 1946) :

71 =



(I/L)”’



or

log n = (pL - pZ)/9



where I is the apparent ionic product and L is the true solubility product

of hydroxylapatite. The factor 1/9 appears because there are 9 ions

involved in the hydroxylapatite formula Ca5( PO, ),OH.

Calculating this index for the solubility of hydroxylapatite in the

presence of calcium carbonate, Ericsson found that it varied with pH

according to the equation:

log n



=



0.44pH - 2.33



Studying the solubility of soil phosphorus in 0.01 M CaCI,, Larsen and

Court (1961) found the relationship between log n. and p H shown in

Fig. 2, from which it may be seen that over the pH range 5.0 to 7.5 the

solubility was consistent with that of impure hydroxylapatite. Above

pH 6.0 the solutions were supersaturated, and below p H 6.0 they were

undersaturated with respect to pure hydroxylapatite. Thus it is not

surprising that Clark and Peech (1955) observed a lower solubility of

phosphorus in acid soil solution than corresponds to the solubility of

hydroxylapatite. Their statement that “At intermediate and low p H values, it is obviously necessary to postulate the existence, in soils, of solid

phosphate phases that are less soluble than the calcium phosphates” can

thus be disputed.

The lack of agreement between the phosphorus concentration of a

soil solution and the solubility of pure hydroxylapatite does not necessarily imply that hydroxylapatite is not determining the phosphorus

concentration.



160



SIGURD LARSEN



A degree of acidity at which all calcium phosphates are so. soluble

that they cannot possibly control the phosphorus concentration will, of

course, eventually be encountered. This may well be from pH 5.0 downward, but it will certainly be true for pH levels below 4.0. At this p H

value the clay fraction will yield significant amounts of aluminum ions

which will then be present in the cation exchange complex and in the

1.0-



0.8-



log n =0.701pH

r = 0.981



-



0.6.

I



0.41

I



-1.oL



FIG.2. The logarithm of hydroxylapatite saturation index ( n ) as a function of

soil pH. (From Larsen and Court, 1961.)



solution. Solubility of aluminum phosphates are likely, in this situation,

to determine the upper limit of the phosphorus concentration of the soil

solution. But more recent evidence has challenged the earlier view that

this limit is governed by the simple solubility product of variscite. Complications of incongruent dissolution and complex formation, analogous

to the chemistry of hydroxylapatite have been pointed out (Taylor and

Gurney, 1962a; Bache, 1963; Raupach, 1963). These are discussed later

in Section IV, B, 2.

2. Adsorption



Even in the absence of phosphorus precipitating ions, phosphorus will

still be removed from solution by adsorption onto the surface of soil



SOIL PHOSPHORUS



161



particles such as clay or calcium carbonate. Adsorption is therefore

another possible mechanism which can determine the phosphorus concentration in solution.

In practice as the adsorption system becomes more saturated by the

addition of phosphorus, the concentration in solution rises and a point

will ultimately be reached when precipitation of a sparingly soluble

phosphorus compound will occur. The solubility of this compound will

then determine the upper limit of the phosphorus concentration; conversely, if the phosphorus concentration is lowered, sparingly soluble

phosphate will dissolve until the adsorption complex has been saturated

to a degree which corresponds to the solubility of the least stable

phosphorus compound present.



3. Chemical Potentials

It is a well known principle of physical chemistry that in any multiphase system at equilibrium, the chemical potentials or partial molar free

energies of all diffusible chemical components are equal. Thus for a

system consisting of a solution phase in equilibrium with a solid phase,

the chemical potential of all components is the same, so that the potential of the solid phase is then easily calculated from activity measurements made in solution.

Schofield (1955) suggested that this approach could be used to

obtain an index of soil phosphate availability. He proposed the “phosphate potential,” the negative chemical potential of monocalcium phosphate (?h pCa pH,PO,) determined in a 0.01 M CaCl, soil extract.

The determination and use of the phosphate potential of soil solutions is,

however, beset with several practical and fundamental difficulties. The

most obvious practical difficulty is the low phosphorus concentration of

the solution which can only be partly overcome by improved analytical

methods. In addition, the pH measurement may give unduly variable

results due to the low buffer capacity of soil extracts.

Less obvious, but very real difficulties are ( a ) lack of equilibrium,

( b ) microbial activity, ( c ) influence of the soil:solution ratio, ( d )

formation of soluble complexes.

a. Lack of equilibrium. Following Schofield and Taylor ( 1955),

Aslyng (1954, 1964) used the chemical potential of calcium hydroxide,

the “lime potential” defined as (pH - %pCa),and the phosphate potential in an attempt to assess the presence and nature of calcium phosphate

compounds in soils. He apparently adopted a short but unspecified period

for equilibration. Similar procedures have been adopted by several other

workers who have recorded their equilibration time. These periods and

other details of the procedures reported are compiled in Table 111.



+



TABLE 111

Methods of Measuring Phosphate Potential

Equilibration time



Method of shaking



Andersen and Mogensen (1962)

Aslyng (1954)

Barrow et al. (1965)

Blakemore (1966)

Chakravarti and Talibudeen

(1962)b

Clark and Peech (1955)

Clark and Peech (1960)



30 min.

Seconds

0-90 hr. (17 hr.)

15 min.

8 days



NSa

By hand

Reciprocating

By hand

End-over-end



Fordham (1963)"



30 sec.-17 hr.



Larsen (1965)d



16 hr.



Author



4 days

4 days



NS

NS



B y hand

End-over-end

Wrist action



Soil: solution ratio

(9. soi1/100 ml.)



Solution conc. (CaC12)

10-2 M



40,SO

40, SO



NS

20

0.5, 5

50

50



10,20

20



M

10-2 M

10-2 M

2 X 10- M KCl



Water, dilute

Water,

2 . 5 x 10-3 M ,

5 x 10-3 M ,

M

10"M

2



x 10-3~,



5



x lo-* M



M,



Larsen and Court (1961)

Larsen and Widdowson (1964)d



16 hr.

16 hr.



NS

End-over-end

Roller

Wrist action



10, 20, 40, 80

2, 4, 8, 16, 32



10-2 M



M



5

z4

2



Moreno et al. (1960)e

Moser et al. (1959)d

Olsen et al. (1960)d

RamaMoorthy and Subramanian

(1960)

Taylor and Gurney (1965)

White (1966)

White and Beckett (1964)

-



1hr.40 days

30 min.

4 days

7 days



30 min.

1 hr.

0-4 hr.



Wrist action



50



NS

NS

NS



40,80

10, 20, 40, 80

10



NS



20,40



End-over-end

By hand

End-over-end



1, 2, 5, 10, 20



10



Water

10-2 M

10-2 M

10-2 M

10-2 M

10-2 M

10” M



~~



Not stated.

b Acid soils only; calculated (+ pM

pH~P04)where M = Ma+ or Fe3+.

Effect of CaCh concentration also studied.

d Partial pressure of COZ controlled.

MCP or DCPD added to suspensions.

a



8



+



B



164



SIGURD LARSEN



The position seems to be that after a comparatively short shaking

period of a few minutes an apparent state of equilibrium is obtained,

followed by a slow increase of the phosphorus concentration in solution.

This slow increase goes on for months, perhaps years, and equilibrium

seems unobtainable.

In this situation a choice must be made. By choosing a short period

of equilibration, a phosphate potential relevant to the most reactive

phosphorus in the solid phase may be achieved. The assumption that

this potential is a measure of the partial molar energy (free energy) of

all the solid-phase phosphorus is not valid, as this is based on the condition that full equilibrium is achieved.

b. Microbial activity. When soil suspensions are shaken for only a

few minutes, microbial activity can probably be ignored, but this is not

the case when the soil is shaken for hours or days. There will then be a

cumulative effect of the microbial activity which, after a short initial

period of adaptation, shows a peak that is particularly marked when

air-dry soil is wetted. This latter effect can be reduced by storing the soil

moist for a prolonged pre-period (White and Beckett, 1964).

There are two main ways in which the microbial activity can influence the amount of phosphorus in the solution and the phosphate potential: ( 1) biological immobilization of phosphorus; ( 2 ) solubilization of

phosphorus by the acidic compounds produced by the microorganisms.

The former effect is generally insignificant, unless the microbial

activity is boosted by addition of organic material. This is because under

normal circumstances the small amount of phosphorus removed from the

soil solution is replaced from the solid phase, so that the effect is to delay,

rather than disturb, the equilibrium.

Solubilization of phosphorus by organic acids from microbial activity

is also a rare occurrence, since such acids are normally quickly decomposed. Accumulation of strong inorganic acids could affect the equilibrium, for example, when oxidizable materials such as sulfur are present.

The most likely cause of solubilization, however, is the accumulation of

carbon dioxide produced by microbial respiration. This difficulty may be

overcome by using a germicide (Barrow et al., 1965), but this has the

disadvantage of possible side effects. Aeration with moist air is the safest

method of preventing accumulation of CO, (Larsen and Widdowson,

1964).

c. Soi1:solution ratio. There have been several attempts to overcome the marked influence that soil: solution ratio has on phosphate

potential. Aslyng ( 1954), for example, extrapolated his experimental results to “zero dilution,” and others have adopted this procedure. This is

not justifiable, however, since both his own data and those of Larsen and



165



SOIL PHOSPHORUS



Court (1960) showed that there is no approach to a limiting value at

zero dilution. White (1966) overcame the effect by allowing the soil

phosphorus to achieve equilibrium by prolonged storage under constant

environmental conditions. Larsen and Widdowson ( 1964) suggested that

an important factor for the soil: solution ratio effect was the accumulation

of CO, during the shaking of the soil suspension in stoppered bottles.

Their results are presented in Fig. 3, from which the marked effect of



ZO



1



2



4



8



16



g. soill50ml. 0.01M CaCI,

FIG.3. Effect of soi1:solution ratio on phosphate potential where COXaccumulation is ( a ) prevented, ( b ) allowed. (From Larsen and Widdowson, 1964.)



CO, accumulation can be seen, an effect which is proportional to the

amount of soil in the suspension. The phosphate potential of this slightly

calcareous soil became completely independent of soil :solution ratio

when CO, accumulation was prevented by aeration. In the light of this

information, the previous work on the effect of soil :solution ratio where

no precaution against CO, accumulation has been taken, should be

reexamined. It is to be expected that different results will be obtained

when the suspension is aerated, except perhaps in the most acid soils

where the biological activity is restricted by the acidity and also where

the carbonic acid formed is largely undissociated.

For example, Salmon (1965) found no effect attributable to CO, in

the two acid soils that he studied. H e ascribed the decrease in potential

which he found with increased soi1:solution ratio to the fact that the soil

was changed more from its original state in attaining equilibrium with

smaller ratios. This would only seem a plausible explanation where the

soil was poorly buffered with respect to phosphorus.

d. Formation of soluble complexes. The calculation of phosphate

potentials is based on the activities of free ions, and no account is taken



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