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Interactions of orthophosphate with iron oxyhydroxide minerals found in soils

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Title:
Interactions of orthophosphate with iron oxyhydroxide minerals found in soils
Added title page title:
Orthophosphate
Added title page title:
Oxyhydroxide
Added title page title:
Iron oxyhydroxide minerals found in soils
Creator:
Yekini, Bourahim, 1948-
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English
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xiii, 123 leaves : ill. ; 28 cm.

Subjects

Subjects / Keywords:
Adsorption ( jstor )
Desorption ( jstor )
Electrolytes ( jstor )
Goethite ( jstor )
Ions ( jstor )
Isotherms ( jstor )
pH ( jstor )
Phosphates ( jstor )
Phosphorus ( jstor )
Soils ( jstor )
Dissertations, Academic -- Soil Science -- UF
Soil Science thesis Ph. D
Soil absorption and adsorption ( lcsh )
Soil mineralogy ( lcsh )
Soils -- Composition ( lcsh )
Genre:
bibliography ( marcgt )
non-fiction ( marcgt )

Notes

Thesis:
Thesis--University of Florida.
Bibliography:
Bibliography: leaves 118-122.
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Bourahim Yekini.

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INTERACTIONS OF ORTHOPHOSPHATE WITH
IRON OXYHYDROXIDE MINERALS FOUND IN SOILS









BY

BOURAHIM YEKINI


A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQM]IEMENTS FOR THE DFGREE OF DOCTOR OF PHILOSOPHY




UNIVERSITY OF FLORIDA


1980































To

The African People













ACKNOWLEDGEMENTS


The author wishes to express his sincere appreciation to Dr. John G. A. Fiskell, chairman of the supervisory committee, for his guidance and assistance throughout the entire course of this study, and for his valuable suggestions and excellent assistance in the preparation of this manuscript. The author is pleased to extend his sincere acknowledgements to Dr. J. J. Street and Dr. V. E. Berkheiser for their participation in the supervisory committee and constructive criticism of this manuscript.

Appreciations are also extended to Dr. W. K. Robertson, Dr. R. C. Stoufer and Dr. T. L. Yuan for various assistances.

Special recognition is expressed to Dr. Charles F. Eno, chairman, Soil Science Department, Dr. B. G. Volk, and Dr. D. F. Rothwell.

Sincere appreciation is expressed to the AfricanAmerican Institute which sponsored this study.

Appreciations are also extended to Carol Giles for her excellence in typing.

The author wishes to express the deepest gratitude to his mother Mariama, his wife Olga and daughter Maryam, and his brothers and sisters for their moral support.


iii














TABLE OF CONTENTS


Page

ACKNOWLEDGEMENTS ........... .......................iii

ABSTRACT .......... ........................... . xi

INTRODUCTION ............. ......................... 1

CHAPTER

REVIEW OF SOME MODELS DESCRIBING FACTORS
AFFECTING NUTRIENT AVAILABILITY ..... .......... 4
Nutrient Potential ....... ............... 4
Capacity and Intensity Relationship ... ...... 7 Energy of Adsorption ....... ............. 8
Some Adsorption Isotherm Models .... ........ 10
Effect of Surface Heterogeneity
on Adsorption-Desorption .. ........... . 12
Generalization of Adsorption
Isotherms ............................ 14
Soil Phosphorus Reaction Mechamisms ..... . 16 Specific Adsorption .... .............. . 19
Infrared Study of Phosphate Specific
Adsorption ...... .................. .. 21

II KINETICS OF ADSORPTION AND DESORPTION ....... . 26 Adsorption ........ ................ 26
Desorption ...... ................. 27
Adsorption and Desorption Relationships . . .. 27 A Kinetic Model for Adsorption-Desorption . . . 28

III THERMODYNAMICS OF ADSORPTIONDESORPTION REACTION .... ............... . 33
The Surface Charge .... .............. . 33
Zero Point of Charge (ZPC) .. .......... 35
Thermodynamics of Adsorption . ......... . 36
Free Energy as a Function of Distance .... . 36 Free Energy for Irreversible Fixation ..... . 38
Relationship of Surface Charge to
Surface Potential .... .............. . 39
Surface Tension and Specific Adsorption . . . . 40










TABLE OF CONTENTS (Continued) CHAPTER Page


IV MATERIALS AND METHODS ...............
Geothite Preparation .............
Some Factors Affecting the Kinetics
of Adsorption-Desorption .... ............
Surface Charge as Affected by Phosphate Adsorption ...... ................
Soil pH, Iron Oxides, and Extractable P . ...

V RESULTS AND DISCUSSION ...... ................
Goethite and Phosphated Goethite Study
by Infrared Spectroscopy ..... ...........
Goethite Structure Identification by
Infrared ........ ...................
Factors Affecting Goethite
Crystallization ............. . .
Phosphate Adsorption and Desorption
Studies . . . . . . . . . . . . . . . . . . .
Effect of the Supporting Electrolyte .....
Time of Reaction Effects ...........
Effects of pH on P Adsorption and
Desorption ........ ..................
Effects of P Adsorption on Surface
Charge of Goethite ...... ...............

VI SUMMARY AND CONCLUSION ....... ...............

LITERATURE CITED .......... ........................

BIOGRAPHICAL SKETCH .....................


99 108














LIST OF TABLES


Table Page

1 Soil pH, iron oxide, and extractable P ... ...... 47

2 Type of salt effect on the P adsorption
on gothite at 5 mg P/ml .... .............. ... 79

3 Effects of three electrolyte salts on the
phosphate sorption maximum and sorption
energy constant for goethite ... ............ ... 81

14 Phosphate desorption from goethite by different anions ....... .................. ... 85

5 Effect of reaction time on phosphate
adsorption maximum and sorption energy
constant for the goethite system .. .......... ... 85

6 The logarithm of (a) equilibrium P concentration as a function of the logarithm
of (b) the amount P adsorbed ... ............ ... 85

7 Effects of P adsorption on the equilibrium solution pH after 24 hours of
adsorption reaction ...... ............... . .101














LIST OF FIGURES


Figure Page

1 Schematic representation of solidinterface-solution system .... ............ . 29

2 Infrared bands of (A) goethite; (B)
phosphate added to goethite at the end
of ageing; and (C) phosphate added to
goethite at beginning of ageing. Curves
A, B and C are for 220C and curves A',
B' and C' are for 550C ....... ............... 50

3 Infrared bands of (A) goethite; (B) phosphate added to goethite at end of
ageing; and (C) phosphate added to
goethite at beginning of ageing. Curves A, B and C are for 220C and curves A', B'
and C' are for 550C ........ ................ 51

4 Infrared bands of (A) goethite; (B) phosphate added to goethite at end of
ageing; and (C) phosphate added to
goethite at beginning of ageing. Curves A, B and C are for 220C and curves A', B'
and C' are for 550C ........ ................. 54

5 Infrared bands of goethite and lepidocrocite. (After Farmer and Palmieri 1975) ........ .. 55

6 Infrared bands of (A) goethite; (B) phosphate added to goethite at end ageing; and (C) phosphate added to goethite at
beginning ageing. Curves A, B and C are for 220C and curves A', B' and C' are for
550C ............ ........................ 56

7 Infrared bands of (A) Fe hydroxide material; (B) phosphate added to Fe hydroxide material at end ageing; and (C) phosphate added to Fe Hydroxide material at beginning ageing. Curves A, B and C are for
220C and curves A', B' and C' are for 551C ..... . 57


vii










LIST OF FIGURES
(Continued)


Figure Page

8 X-ray diffraction patterns of (A) goethite, (B) phosphated goethite at end of ageing, and (C) phosphated goethite
at beginning of ageing .... .............. .... 59

9 Infrared bands of phosphated goethite at beginning of ageing for suspensions
of various P/Fe values at OH/Fe = 6 . . ........ 62

10 Infrared bands of phosphated goethite at
beginning of ageing for various P/Fe ratios at OH/Fe = 3.0 ..... ................ ... 63

11 Infrared bands of phosphated goethite
at beginning of ageing for various P/Fe
ratios at OH/Fe = 1.5 ....... ............... 64

12 Infrared bands of 1) goethite digested in
D20 and 2) phosphated goethite digested
in D 20 ....... ....................... .... 66

13 Infrared band of phosphated goethite
after desorption by (1) 0.1 N KCl, (2) H 20,
(3) 0.1 N KNO3, and (4) 0.1 N Na2SO4 ........ 67

14 The 001 face of goethite lattice. (After
Bragg and Claringbul 1965) ... ............. ... 69

15 The 001 face of phosphated goethite . ........ . 69

16 Adsorption isotherms as affected by the
reaction times for goethite-solution
systems ........ ...................... ... 71

17 Adsorption isotherm for the Kenya soil
after 24 hours of reaction time . ........... ... 72

18 Adsorption isotherm for the Georgia soil
after a reaction time of 24 hours . ......... ... 73

19 Phosphate adsorption isotherm for the Colorado soil after 24 hours of reaction time .. ..... 74


viii










LIST OF FIGURES
(Continued)


Fjr Page

20 Effects of the initial P concentration on the phosphate adsorption by goethitesolution (1 g/1000 ml) .... ............... ... 76

21 Effects of the type of supporting electrolyte on P adsorption on goethite
suspension (1 g/1000 ml) .... .............. ... 78

22 Effects of the supporting electrolyte
concentration on the kinetics of P adsorption by the goethite-solution system
(1 g/lO00 ml) ...... ................ .... 84

23 Phosphate desorption in 10 hours from
goethite by three salt solutions at various equilibrium P concentration .. .......... ... 87

24 Transformed Langmuir equations for phosphate sorption by goethite as affected by reaction times. a is equilibrium concentration and b amount adsorbed .. ........... ... 80

25 Change in the equilibrium P concentration
and P sorption as affected by the reaction
time .... ......... ................... ... 89

26 Change in the equilibrium P concentration
for three soils as affected by the reaction
times ........ ....................... ... 90

27 Effect of reaction time and equilibrium P
concentration on P desorbed by 0.01 N CaCl2
from Kenya soil ...... ................. ... 95

28 Effect of reaction time and equilibrium P
concentration on P desorbed by 0.01 N CaCl2
from Georgia and Colorado soils .. .......... . 97

29 Effect on P desorbed by 0.5 N NH F as
affected by the adsorption reaction time
and equilibrium P concentration .. .......... . 98










LIST OF FIGURE
(Continued)


Figure Page

30 Logarithmic plot of equilibrium P concentration change with time (t) relative to
equilibrium at 18 hours (t ) for phosphated
goethite . .....................93

31 Effects of change in equilibrium solution
pH on P adsorption and 0.01 N CaCl desorption of P in goethite system. Initial P
concentration is 5 pg P/ml .... ............. . 100

32 Effect of phosphate adsorption time on
phosphate desorption by 0.01 N CaCl2 at
pH 2 and pH 10 ....... ................... . 104

33 Effects of solution pH on the amount of
P desorbed from goethite .... .............. . 106

34 Effects of pH and concentration of supporting electrolyte on phosphate desorption 107

35 Potentiometric titration of goethite. Note
zero point of charge occurs at pH 5.8 . ....... . 109

36 Potentiometric titration of phosphated
goethite. Note zero point of charge occurs
at pH 5.2 ............ .................. 110









Abstract of Dissertation Presented to the Graduate Council of the University of Florida in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy

INTERACTIONS OF ORTHOPHOSPHATE WITH
IRON OXYHYDROXIDE MINERALS FOUND IN SOILS By

Bourahim Yekini

June 1980

Chairman: Dr. John G. A. Fiskell Major Department: Soil Science

The mechanisms, kinetics, and reversibility of orthophosphate adsorption on synthetic goethite and soils were investigated. Goethite was prepared by ageing of the precipitate which appears after mixing FeCl2 and Na0H solutions. Both the increase in Fe/OH ratio in the suspension and in ageing temperature favored a higher degree of goethite crystallization within a short period of time. With either OH/Fe = 6 in the suspension at room temperature or OH/Fe = 3 at 55C and ageing for one week, good goethite yield was obtained. Infrared bands characteristic of goethite were at 3200 cm-1 (OH streching) and at 890 cm-1 and 790 cm-1 (both are Fe-OH bending vibration). The goethite structure was confirmed by X-ray diffraction intensities of the 4.19 R and 2.70 R peaks. The presence of phosphate at the beginning of goethite ageing weakened the goethite structure. When F/Fe = 3.2 (regardless of OH/Fe ratio) the bond (Fe)- O-P vibration at 1000 cm-I was predominant, thus preventing the formation









of Fe-OH bond even after ageing of a suspension. When the addition of phosphate was made at the end of the goethite ageing, absorption of phosphate was determined as surface binding through binuclear bridging of the HPO42- ion which was identified by the presence of bands at 1120, 1085, and 1030 cm . The phosphate was assumed to penetrate the goethite structure whenever the vibrational band at 1000 cm1 was present.

Using synthetic goethite as supporting medium, it was found that the Langmuir and Freundlich equations could be used to describe the relation between the amount of P adsorbed and that remaining in the equilibrium solution for particular periods of reaction time.

A kinetic model was proposed to describe the change

with time of each of the following ion forms: 1) the free ions in solution, 2) the physically adsorbed, 3) the reversible chemically adsorbed ions, and 4) the irreversible chemically adsorbed ions. The amount of phosphate adsorbed increased with the increase in initial P concentration and the adsorption reaction time. The time required for the goethite phosphate system to reach an equilibrium state increased with the initial P concentration. Generally, the equilibrium state was reached after 18 hours for goethitesolution and 22 hours for soil-solution systems. The adsorption on goethite was increased by both multivalent cations


xii









and the concentration of the supporting electrolyte. The potential binding energy constant increased (from 0.70 to

4.60 ml/ pg P) as the reaction time increased, but decreased as the cation valence of the electrolyte was increased.

At the same initial P concentration, the amount of P

adsorbed decreased almost linearly with change in pH according to the relation pg P/g goethite = -446 pH + 5680. The amount of P desorbed over a wide range pH or supporting electrolyte remained nearly constant when the desorbing time was greater than or equal to 6 hours. The P desorbed from goethite and/or soils increased both with the initial equilibrium P concentration employed and the reaction time. At the same initial P adsorbed on goethite, the amount of P desorbed from goethite decreased when the pH was below 5.5 and increased when the pH was greater than 6. Phosphate adsorption on goethite was found to induce a net increase of negative charge so that the zero point of charge declined from pH 5.8 to 5.2.


xiii













INTRODUCTION


Increase in urbanization makes it necessary to increase both the quantity and the quality of agricultural production in the tropical region. The implementation of an adequate agricultural policy involves improvements of the actual level of technology. This implies a clear understanding that the macroscopic relation be sought between the major factors producing food for human population through proper managements of soils, plants, animals, and insects under diverse climates and social activities. Modeling of the agricultural system is a useful tool for understanding a united approach for all major factors governing the system. With a suitable model, it should be possible to envisage needed change in a particular factor in order to give results as close as possible to a reasonable expectation. Only some microscopic relationships within the soils will retain our attention in this study. In a soil-solution system, the dominant soil phenomena taking place simultaneously are mass transport, diffusion adsorption-desorption, precipitationdissolution, and microbial immobilization and mineralization. The dominance of each soil phenomenon depends on the soil structure and texture, organic matter content, water content, temperature, and the ionic suite (types and concentrations).









Upon fertilization, tropical soils may not react in the same way as do the temperate soils because of their different mineralogical, and chemical properties. Among the major elements, phosphorus is, after nitrogen, the most deficient nutrient and its availability is strongly dependent on the mineralogical composition of the soil (type of clay, and metal oxides). The low availability of phosphate to tropical plants is due both to the high soil phosphate fixing capacity and to precipitation of fertilizer phosphorus from the soil solution. Not all forms of phosphorus bound on the oxide surfaces are held with the same degree of strength. With time, as phosphorus is depleted from the solution, some phosphate may be replenished by release from the solid phase. The magnitude of this release, with respect to time, is dependent on soil characteristics. The initial concentration in solution is not a sufficient measure of phosphorus availability. In tropical soils, iron oxides are quite important in determinig the solid phase capacity to fix and supply phosphorus to the soil solution. How some factors, such as OH/Fe ratio, temperature, and phosphate concentration affect the formation of goethite are herein investigated. Infrared spectroscopy will be used to identify the effects of soluble phosphate on goethite structure and the nature of phosphate binding on goethite. Synthetic goethite as well as three soils will be used to investigate some





3


factors affecting the time dependance and the degree of reversibility of orthophosphate adsorption.













CHAPTER I
REVIEW OF SOME MODELS DESCRIBING FACTORS
AFFECTING NUTRIENT AVAILABILITY


Nutrient Potential

The concept that availability of a nutrient can be described by its potential activity was first brought to general notice by Schofield (1955). His concept was that in comparing two soils of different water holding capacity which contain the same amount of available water, the one having the lower capacity has a lower potential and the water is easier to extract than from the soil of higher holding capacity. Their pF value defines the energy with which water is held on the soil particle surface. The potential given by the pF value provides the basis for quantitative evaluation of the availability of water. As water is taken out of a soil, it is replaced by lateral movement of groundwater. Similarly, as a nutrient is taken from the soil solution, it is replaced by other ions, by desorption, or by diffusion as well as mass flow. Eventually, it is not the amount of a nutrient in a soil that primarily controls the uptake of that nutrient by the plant, but the work required to withdraw it from the solution. This work may be related to Gibb's free energy (Gi) if the uptake is mainly from the soil solution. To derive an expression for





5



the nutrient potential, three postulates have been advanced by different workers.


Postulate 1

Where the solution is at constant temperature and pressure, and when the system is at equilibrium, Gi is uniform throughout the soil solution so that

G. = G0 + RTLn a. + zF. (f)
1 1 I i
where G? is the standard molar free energy of an ion i in
1
the solution relative to an arbitrarily zero of the electric
f
potential, a. is the activity of this ion in solution, z is
1
the valence, F is the Faraday constant, and T is the
1
electric field effects.


Postulate 2

When a soil is in equilibrium with a solution, the

electro-chemical potential of the ion is constant throughout the system (soil-solution). Then, the partial molar free energy of any ion in the soil complex can be determined by analyzing the solution. This free energy of the ion in solution is assumed to be constant even if the solution is separated from the solid phase by centrifugation. As a result, if the ion is taken out of the field influence, then zFi' = 0 where z, F, and 4' are as described in the previous section. Then the relationship in Eq. (1) can be reduced to










G. = G0 + RTLn a. (2)


Postulate 3

This postulate is that potentials are determined by transfers of ions. In order to maintain the electroneutrality, equal transfer must occur for quantities of positive and negative ions in one direction or, alternatively, transfer of one ion in the first direction occurs as an equivalent quantity of ions of the same sign moves in the opposite direction (Barrow et al. 1965). It is also assumed that divalent cations (Ca and Mg) dominate most soil surfaces as exchangeable cations and in the solution, except for saline soils. For a situation in which transfer of Ca2+ is in one direction and K+ in the other, the net change in free energy is
AG AGa0 + RT(ln aK - Ln a (3)
Ca,K Ca,K KCa

For phosphorus, over the range of pH which exists in soil, two forms of the orthophosphate ions are in equilibrium; thus

H + 20 7; HP0H042- + H0 + (4)
2 4 2pK = 7.2 43

Assuming that plants absorb mainly H PO ion and not HPO22+4
then a cation such as Ca2+ must accompany the H 2P04' The equation for the transfer of both ions from the equilibrium solution is









AGCa(ll2P04)2 = -2.3 RT(l/2pCa + pH2P04) (5) A common use of the relationship is the Schofield's phosphorus potential, where I is given by I = 1/2 Ca2 + pH2PO 4'


Capacity and Intensity Relationship

Phosphorus ions pass from one phase to another as a

result of chemical potential differences between the phases. In general, any substance tends to pass from a region of higher chemical potential to a lower one. The ability of the solid phase to supply phosphorus to the solution phase, as it is depleted, can be termed the capacity. For any ion such as phosphorus, at equilibrium, the phosphorus buffering capacity (PBC) is a soil characteristic where PBC is dAP/dI. The term with I = I/2pCa 2+ + pH2PO4 and AP is the gain or loss in phosphorus by the soil. The relationship between capacity and intensity factors is expressed as Q/I curves which are generally composed of two parts with linear and curvilinear portions. In the Q/I relationship for potassium, the curved portion is attributed to a number of sites which have specific affinity for potassium at low concentration and the linear part is associated to non-specific sites, Beckett (1971). He also believed that, in general, the curved part of the Q/1 relation can be represented by a Langmuir adsorption isotherm whereas the linear part fits the Gapon relationship.









Beckett (1971) affirmed that the non-linear part of the curve represents a certain region where there is a definite and limited number of sites on the exchangeable surfaces and exhibits a selective binding power for the adsorbing ion: The Langmuir adsorption isotherm could be used to describe such a situation. In addition, if Eq. (3) is recalled, the activity ratio (AR) is K+/(Ca)I/2 of the ions in solution which is the product of the concentration of the exchangeable ions and the Gapon constant (kG).


Energy of Adsorption

Partition Function

To be effective, the collision between molecules and

the collision surface must provide a certain minimum amount of energy called the activation energy. Since the activated complex is a transitory species, the equilibrium constant cannot be measured experimentally. However, the partition function arises from the quantum theory that a molecule can exist only in states with definite energy limits. From Boltzman distribution, it is recognized that

N A = N oA/g oAgiA exp(-E iA/kT) (6) where NA is the total number of molecules, NoA is the number of molecules at zero energy, E. is the activation energy at the state i, and gi is a constant. In this case, also quantity Q = IgiAexp(-EiA/kT) which is the partition function.









Electric Field Effect

The following treatment was first derived for solid-gas interaction and is known as Thin's theory reported by Brunauer (1943). The strength of the electric field surrounding the adsorbent is given by E = e1/er0, where e1 is the charge distribution of the absorbent, E is the dielectric constant of the gas, and r is the distance from the surface.

The force (F) acting on an induced dipole is

F = ke v/4 nr2v+I
1

where v is a constant and k = E-I/E. The adsorption potential is the force multiplied by the distance through which its acts (F.Ar).


Heat of Adsorption

Assuming no work due to phase change is done during the adsorption-desorption, the molar energy of the ion in free solution (u A) and the energy of the ion in the adsorbed phase (u B) are related to the loss of ion from the solution due to adsorption as shown by
MH = b(uA - u B) (7)


where b is the number of mole adsorbed, and AH is the integral heat of adsorption.


Differential Heat of Adsorption

Assuming. that the adsorption process is reversible All

can be obtained from data at two temperatures (T1 and T2) by








using the Clausius-Clapeyron equation for a constant surface coverage, namely


ln _ AH 1(8) L alj R T T21

where a1 and a2 are equilibrium concentrations at temperature T1 and T2, respectively, and R is the molar gas constant.


Some Adsorption Isotherm Models Freundlich Equation

The Freundlich equation was first introduced in an

empirical form (Bach and Williams 1971). It assumed that the energy of adsorption decreases exponentially with increasing saturation of the surface. The equation is

b/m = ka1/n (9) where b/m is the amount of phosphorus adsorbed per unit weight of soil at the equilibrium concentration termed a, and where k and n are constants. The logarithmic form of Eq. (9) is

log(b/m) = log k + 1/n log a (10)

The plot of log(b/m) versus log a should yield a

straight line. This equation is valid only in a limited range of concentrations. By taking into account the initially exchangeable phosphorus and the ability of the soil to lose or gain phosphorus (AP) during the equilibrium reaction, the following modification was introduced

AP Aa = ka1/n- e (11)









where Aa is the amount of phosphorus gained or lost, e is the phosphorus initially present, and n is a constant. Modified Freundlich Equation

In a colloid-solution system, the adsorption is timedependent. Kuo and Lotse (1972) introduced a time factor l-e-k2t after assuming that the constant k2 is small. They found that

b/m = ka t1/n (12)
0
where a�0 is the initial phosphorus concentration. Their study indicated that the rate constant increased with an increase in concentration.


Langmuir Equation

The derivation of the Langmuir equation is based on

the following assumptions: 1) the energy of adsorption is constant and independent of the degree of coverage, 2) there is not interaction between adjacent adsorbed molecules on an homogenous surface. If the system is in dynamic equilibrium, this results when the rate of adsorption is equal to the rate of molecules escaping from sorption surface, so that

k1e = k2 P(l - 6) (13) where k and k, are the rate constants of adsorption and
1 4
desorption, P is the gas pressure, and 6 is the fraction of the surface coverage by the gas. By rearrangement, the relationship is









ksP (14a) and by analogy

ksa (14b) for soil-solution system, where a is the ion concentration in solution. The above equation can be put in the linear form as follows

a/b/m a/s + 1/ks (14c) where k is the adsorption energy constant and s is the adsorption maximum. Different workers have observed that there is not a linearity between a/b/m and a. This may be caused by at least two types of sites of adsorption having different energies of adsorption. To fit this criterion, the Langmuir equation was modified to be
klsla ksa
1 1 + 22(14d)
1 + kla 1 + k2a

where kl, sl are constants for region 1, and k2, s2 are constants for region 2 (Syers et al. 1973).


Effect of Surface Heterogeneity on Adsorption-Desorption

Langmuir and Freundlich equations can be obtained from Toth's equation (Jossen et al. 1978), for homogeneous surface. Toth's equation is obtained by integrating the equation

= dlna - 1aa (15) Where B = 0, the Freundlich equation is obtained. Where B = 1, the Langmuir type of equation is










b = a (16) where 6 = d exp(-E/RT), a is the equilibrium concentration, b the amount adsorbed and bo is the maximum adsorption.

Another factor is that the energy of activation (E) is constant for a homogeneous surface but varies on a heterogeneous surface with the degree of coverage (Jossen et al. 1978). Assuming that the free energy of activation for adsorption varies linearly with coverage of the surface, Lingstom et al. (1970) proposed the following model

A + S AS;- AS (17)

2

where A is the adsorbate, S is the surface, and k and k2 are the rate constants. They deduced that the rate of adsorption process is

de/dt = k1(l-)(1-1/20)e-be + k2[(1-i/20)e- b(2-1)_/202 e b

(18)

where e = b/b.

Another useful relationship is the Elovich equation where the activation energy is a linear function of the amount adsorbed

E = E + ab (19a) Aharoni and Ungarish (1977) modified this equation by introducing the fact that the heterogeneous surface is comprised of a large number of homogeneous regions having unequal number of adsorption sites.









Et =E RTln(gyt/nE + y) (19b) where yt = db/dE, g and y are constant, and nE is the number of sites of energy Et, where Et is the activation energy characteristic of a region. They also assumed that, at any moment, adsorption takes place preferentially on the region that has the lowest activation energy at that moment. Considering the case where equilibrium is attained at a region when yt = Yeq, then it follows that y = k cexp(E/RT). The overall rate of adsorption is given by

db/dt = koNtexp(-Et/RT) and Nt = fnEdE (20) where Nt is the number of sites with an energy Et at time t and nE is the number of sites with an activation energy between E and E + dE.


Generalization of Adsorption Isotherms

Freundlich and Langmuir equations can be extended for cases of competitive adsorption (Jaroniec and Toth 1976; Digiand et al. 1978). In this case, the Freundlich type is given by

b= k(Ea.) (21)
i.

and the Langmuir type by

ka /(k + k a + k a + k.a.) (22)
1 1 0 11121

which is a partial isotherm of a1 relative to the total ions present. A binary system can be simplified if the following









assumptions are made, 1) the surface is formed by a collection of regions each being characterized by a binding energy

(E), 2) the number of surface atoms Ni with energy E follow the Boltzman distribution, which is described by

Ni. = N0exp(-E/RT) (23a) and 3) all adsorbed molecules stay on the surface for the same average length of time 6 so that

6= 60exp(E/RT) (23b) During the time 6, the fluctuation of the number of adsorbed and desorbed molecules compensate between themselves randomly (Vlad and Segal 1979). They considered that the adsorption energy is an increasing function of the extra energy c. From Mclaurin development, the extra energy can be expressed as
= Zn (E - E )n (24)

n

where Em is the smallest adsorption energy corresponding to the value of zero energy and n is generally integers of 1 or

2.

The energy distribution X(e) is expressed by Van Dongen approximation as

X(E) = exp(E a.E) i = 0, 1, or 2. (25a)
i

or generally as

N n-l N
X(E) = l/kT I nc (E - E )n exp[-l/kTZEa n(E - E )n] (25b)
n n n n








The general isotherm is expressed as

b(a,T) = f b1(a,E) X() de (26) where b(a,T) is the fraction of the total surface covered at (a,e), b1(a,c) is the local isotherm which may be analogous to an isotherm equation for an homogeneous surface or region, Q is the range of possible variation of adsorption energy assuming that Q = [0,].


Temperature Dependence of the Rate Constant

According to the Arrhenius equation, it is known that

kads = Aexp(-Ea/RT) (27a) where Ea is the energy of activation, R is the gas constant, and T is the absolute temperature.

From the transition state theory, this relationship is extended to

kads = (kT/h)exp(AS a/R)exp(-AH a/RT) (27b)

= (kT/h)exp(AGa /RT) (27c) where ASa is the entropy of activation, AH = Ea - RT which is the enthalpy of activation, and AGa is the Gibb's free energy of activation.


Soil Phosphorus Reaction Mechanisms

Plant responses to soil phosphorus are a function of the solubility of phosphorus. Any factor altering the solubility of phosphorus will also alter the plant response









(Hemwall 1957). The solubility of phosphorus depends on several factors such as the composition of the soil solution, the pH of the solution, and temperature of the soilsolution system.


Calcareous and Neutral Soils

The adsorption of phosphorus on calcareous surfaces can take place by replacement of water molecules, bicarbonate, and certain other anions or cations. The relative strength of the phosphate ion adsorption depends on the solubility of the compound at the calcium surface (Kuo and Lotse 1972).

The precipitation of phosphorus may be due to the

formation of a whole series of insoluble calcium phosphates. Some of these with associated solubility products expressed as the pK value are reported by Lindsay and Moreno (1960).



Compounds Chemical formula pK

Calcium phosphate anhydride CaHPO4 6.66 Calcium phosphate dihydrate CaHP04. 2H 0 6.56 Octocalcium phosphate Ca4H( P0 4))3.3H20 46.91 Hydroxyapatite Ca 0(P0 4) 6.(OH)2 113.70 Fluorapatite Ca10 (PO4)6.F2 118.40


In calcareous soil, hydroxyapatite and fluorapatite are the major phosphorus compounds, whereas in neutral soil, octocalcium phosphate becomes important and in moderately acidic soils dicalcium phosphate may occur.









Acid Soils

Iron, aluminum and pH are the main factors controlling phosphorus solubilities: In acid soils, the solubility of phosphorus increases with decreasing free iron and aluminum activities and with increasing pH. The form of phosphate sorbed by acidic soils was recovered (90%) as iron and aluminum phosphates (Ghani and Islam, 1946). Yuan et al. (1960) reported that up to 80% of the added phosphorus to acid sandy soils was present as aluminum phosphate and 10% as iron phosphate, but when the reaction temperature was increased more iron phosphate was formed.
3+
After reaction of phosphorus with soluble Al , a

microscopic examination showed an hexagonally shaped crystal in which the interplanar spacing was similar to those of palmerite (Haseman et al. 1950). In general, some of the main forms in which phosphorus can precipitate with iron and aluminum are given below.



Compounds Chemical formula pK
Variscite AlPO 42H 20 30.5 Strengite FePO 4. 2H20 34.3 Taranakite (K,NH 4)3H6Al5(P04)8.18H20 176.0 Palmerite HK2 Al 2(P0O)3H20



Since Al and Fe atoms are part of the surface colloids, they react with phosphates. Whether the process is









precipitation or adsorption depends on the size of the metal polymer and the pH of the phosphate, and its concentration. In a moderately acidic medium, with a high phosphorus concentration, the reaction process may be typically precipitation, resembling that reported by l.su (1965), [A16(OH)I2 6+ + 6 H2PO_ _ -- Al6(OH)2 (H2PO4)6 (28)


Specific Adsorption


General Description

In a solid-solution system, adsorption of molecules (or ions) occurs when there is a change in phase from the free state (in solution) to the bound state at the interface. Ions deposited at the surface likely orient to form the Stern layer as some ions may approach closely to the surface structure. In this case, the ions are said to be specifically adsorbed on the surface.

Non-specifically adsorbed ions are either in the diffuse Gouy-Chapman region separated from the solid surface by at least one molecule, or electrostatically bound to that surface. Specifically adsorbed ions are in the coordination shell of the surface atoms and are maintained there through chemical binding (covalent or coordinate binding). Since specific adsorption of cations or anions occurs even when the surface possesses a net positive or negative charge, respectively, there must be an electrostatic contribution due to polarization of the ion or molecule.








Both chemical and electrostatic attractions contribute to the energy of adsorption. The magnitute of this energy of adsorption determines the degree of reversibility and the amount adsorbed. Since the reaction is pf1 dependent, at any pH, there is a maximum adsorption of anions and when these maxima are plotted against pH, the curve is termed the adsorption envelope which is described by Hingston et al. (1967) as

A = 2V1(i - ) = 2V( KHJ 2 (29 V1 (K + EH]) (9
where A is the amount of the ion adsorbed per unit weight of adsorption, a is the degree of dissocation of acid anion, K is the dissociation constant of the most highly charged anion that is adsorbed, and V is the amount of ion adsorbed at the maximum level.

Why some of the adsorbed anions such as phosphate may be irreversibly held is explained by certain postulated mechanisms of phosphate adsorption. The removal of phosphate from free solution is assumed to proceed through the replacement of coordinated H20 groups and/or some OH ions,


OH

H OH O-P=O
2 1 / I
OH 2 OH + OH (30)
\OH OH "-2OH
OH


where N is a metal (Fe or Al), and bonded by a coordination link.









Infra-red Study of Phosphate Specific Adsorption Theory of Infra-Red Spectroscopy

If a molecule is placed in an electromagnetic field, a transfer of energy from the field to the molecule occurs when Bohr's frequency condition is satisfied: AE = hv, where AE is the difference in energy between two quantized states, h is Plank's constant, and v is the frequency of the light wave. Pure rotational vibrations are usually observed in the microwave vibrational spectra in the infrared whereas electronic spectra are in the visible and ultraviolet. The infrared spectra originate in the transition state existing between two vibrational levels of the molecules. From a classical point of view, a vibration is active in the infrared spectrum if the dipole moment of the molecule is changed during the vibration. The dipole moment Q is related to the strength of the electric field by Q = cE and Q is a vector whose direction is the line between the center of gravity of the protons and electrons. Let a diatomic molecule be represented by two masses m1 and m2 moving along the molecule's axis with displacements of X and X2 respectively. The displacements of the two atoms are induced by forces which can be obtained through Hooke's and Newton's law.

d2X1 d2X2

K(X - X- XI)= m (31)
2 1 t2 and 2 1 2 dt 2








The solutions of these equations of motion are:

X1 = Acos(2nvt + (x) and X2 = Acos(27vt + a) (31b) After differentiating and substituting back in Eq. (31), it can be found when solving for v (Nakamoto 1978; Colthup et al. 1975)

V 1 m l + 1 (32)
2 7T m 1 2

where k is the force constant and v is the frequency of vibration.

The experimental observations of frequency of vibration of a crystal are due to bond vibrations of atoms within the unit cell. Such vibration may help us to determine the nature of the atoms composing the unit cell by comparison to known polyatomic vibrations. Because of the interactions, the symmetry of a molecule is generally lower in the crystalline state than in the isolated state. The isolated P034
ion is of tetrahedral form (Nakamoto 1978):


0 %.0 0 0 P P P P
0 Oo 0 0

V1 V2 V3 V4


The above tetrahedral structures predict two infrared active fundamentals, one is stretching (v3) and the other (v) is bending.






23



Infrared Identification of the Forms of P Adsorption

From their studies, Hingston et al. (1974) proposed the reaction


OH2

Fe O0 0 0-P 0 O


'-OH 2


= 0


Reversibly adsorbed P


OH2

Fe 0



OH2


Irreversibly adsorbed P


Parfitt et al. (1975) identified a form of phosphorus

which can be considered to be completely within the goethite structure, since the phosphate is bound with two Fe atoms and is H-bonded to a third Fe.


Fe - O---HO


0/


0 - Fe


P


0 - Fe


Fe - OH---O 0 - Fe


0/
0 0 -Fe


Such a chemical binding of phosphate on the goethite surface is said to be specifically adsorbed, and can be considered as the formation of a new solid phase or growth of the solid phase.


Colloid or Soil Surface Effects

Substances dissolved in soil solution can move by

molecular or ionic diffusion resulting from a concentration


-i


0


+ OH


(33)


(34)









gradient within the solution, or by mass flow of the soil solution. This is complemented by sorption of the ions onto the soil surface, or by a combination of these factors. The above phenomena occur simultaneously and are described by the following equation of de Camargo et al. (1979):
2
da A d a B da C db dT (35)
dt- dx dt dx

where a is the ion concentration in the soil solution, b is the amount of the ion adsorbed, x is the distance of the ion from the electric potential effect due to the surface charge on the charged solute. In Eq. (35), the constants are A which is a coefficient combining diffusion and hydrodynamic dispersion (cm 3/hour), B is the average pore density derived from the volumetric water content and C is the average pore water velocity, and D is the ratio between soil bulk density and volumetric water content.

In a soil system where the terms Bda/dx and DdT/dx

approach zero in a soil column of a semi-infinite length, over a small time t, Eq. (35) can have the following solution given by Lingstrom et al. (1968):


a(x,t) = aoerfc 2 (36) where a is the initial concentration, and K is a constant depending on the free energy of adsorption.








Root Effects

The effect of roots in a soil solution arises because

ions movements occur into and out of the roots, with the net balance being an influx. This uptake is mainly associated with ion transport by diffusion and mass flow and is proportional to the concentration at the root surface. It is assumed that the rate of plant uptake is equal to the rate of loss of the solute, so that the reaction is:

da -Sa (37)


where t is the reaction time, and 6 is a constant. A solution of the above differential equation is given by Baldwin et al. (1973) in the following form:
2
w 2x x)(8
a =aexp wt/(W + 2 ln r (38) where a is the initial concentration in solution, and where
0
w, wl, and w2 are constants, r is the root radiu,, and x is the distance from the root surface. This equation describes the change of the solution concentration at a certain distance from the root surface with time.













CHAPTER II
KINETICS OF ADSORPTION AND DESORPTION


The interfacial region between soil colloid and solution is a center of intensive chemical and physical activities. The type of activities retaining-our attention here is the adsorption and desorption processes of ions.


Adsorption

The adsorption of an ion on the colloid surface can be considered as a second order reaction involving an ion (A) and the sites (S) on the surface:


A +S AS (39) where k is the rate constant of adsorption.

By assuming that the reacting sites on the colloid surface are reacting species equivalent to (A), then

dA = k [A]2 (40a)
dt 1

and rearranging this is

d A = ftk dt (40b)
[A]2 t=O

or can be written as

1 - k1 1t (41)
[A] (A] 0

The plot of 1/[A] versus time t should yield a straight line.









Desorption

The desorption of any adsorbed ion (AS) can be considered as a pseudo-first order reaction:
k
A + S AS (42) where k2 is the rate constant of desorption. By differentiation of Eq. (42), the following can be obtained:
d[Al k2[A] (43)

dt 2

And by integrating the above differential equation, this becomes

[A] = [A]0exp(k2t) (44) where [A] is the initial concentration of the ion specie A, t is the reaction time, and k2 is the rate constant of desorption.


Adsorption and Desorption Relationships

Since the adsorption and desorption phenomena are

simultaneous a combination of the Eq. (41) and Eq. (44) is formed as follows:

k
A + S ---- AS (45)
2

d[Aw/dt = -k 1 [A] 2 0i6)



where A, S, t, kand k 2are described in the precedent









sections. By dividing the Eq. (46) and Eq. (47), and rearranging, it can be obtained:

d[Al k 1 d[A]
d[A]2 k2 [A]

1/[A] - 1/[A] = (k1/k2 )ln([ASI/[AS]� (48) If we set a = 1/[A] - 1/LA]0 and b = [AS], then Eq. (48) can be written as
k2
b = b� exp(T- a) (49)
0 1

where b is the amount initially adsorbed.

The above equations conbining adsorption and desorption phenomena are valid only at equilibrium, there is need to propose another model combining adsorption and desorption but valid at any reaction time.


A Kinetic Model for Adsorption-Desorption Energy Constant Characteristic of the Phase

For a collision to be effective in producing molecular species from reactants, some amount of energy must be available to allow for the necessary bonds to break and be formed. In 1889, Arrhenius suggested that molecules must get into an activated state before they become reactive. In any system, an equilibrium exists between ordinary and active molecules and only the latter are rich enough energetically to undergo reaction









A (A-B)# B (50) where (A-B)# is the spatial configuration of the transition state. The species (A-B)# does not represent the active molecule of Arrhenius. His active or energized species are rather a few reactant molecules having sufficient energy to get into the transition state, (A-B)# , or activated complex condition, but not necessarily having the spatial configuration corresponding to the transition state (Eying and Eying 1963).

Let us consider the following system



11I II IA
II1C I/1 -B A A

/4/I/-//---------/7777777



Figure 1. Schematic representation of solid-interfacesolution system.


where A is the solution phase, B is the interfacial phase, and C is the solid phase. Here A, B, and C are different degrees of energized species of the same ion (or molecule) with respect to the solid surface; these are considered characteristic of solution, interfacial, and solid phase, respectively. The surface sites at which ions take the









form C are at random on the solid surface. It is believed that the sites associated with the lowest energy are the first to be filled by the ions taking the form C.


Formulation of the Model

In the system described above, it can be assumed that the following type of reaction is taking place:

k k
A B C (51)
-1 -2

where A, B, and C are as described previously and subscripts a, b, and c designate the amounts of the ion in the forms A, B, and C per unit volume of the system. The terms kI k2 are the rate constants for forward reaction and reverse reactions, correspondingly are k-1 and k-2.

The rate of a reaction can be expressed by the rate of change of the concentration of the ion species:

da k_ b - k a (52a)
dt -1 1
which in expanded form becomes db = k a - k b + k 2c -k 2b = k a - (k- +k )b + k 2c (52b) dt 1 --2 2 1 -l2and

dc k b -k_ c (52c)
dt2 -2
Since any colloid cannot, within a finite period of time, indefinitely fix an ion in the C energy form, then c must reach a maximum s at equilibrium. In Eq. (52b), it is assumed that k_2 is very low compared to k2.








The simultaneous solution of the differential for Eq. (52a) and Eq. (52b) is

w1t w2t
a = ae and b =Be (53) where


aw1ek = kee1 + k15e (54a) and

ww2e k 1lte (k_ 1+ K2)Be2 (54b)


From the Eq. (54a) and Eq. (54b), the coefficients a and 8 can be determined as a function of the rate constants.

In any soil-solution system, equilibrium is attained when the rate of change of each ion in the different phase becomes nil, so that it is then

da db dc dt dt dt

such that k2b - k_2c = 0 and c=s, then the term B is deduced as

k_2 -w2to
= -k-2-0(55)


Similarly, if k_1b - k1a = 0, then substituting in the 6 value, a is found to be

k_1 k_2 1 to
k k e (56)


In the initial a and b expressions, shown in Eq. (53), it can be shown that








k_ k_ 2 wl(t-t)
a =s - T e (57a) and

k_2 w2(t-to)
b = s 7- e (57b)


The Eq. (52c) can be expressed in the following as a first order linear differential equation:

de w~t
dt + k_- c = 6ke2 (58) The solution to this differential equation to satisfy the condition where c = 0 at t = o can be written in the form


C k2 ew2 - e k2t] (59a) If we assume that k 2 is negligible, c can be further simplified

k2 w2t
c= k e (59b)
wnw2
where wI1 and w 2 are constants.













CHAPTER III
THERMODYNAMICS OF ADSORPTION-DESORPTION REACTION


At the microregion in the colloid or soil solution

extending from the surface to the outer limit of the first adsorbed layer, it is assumed that the electrical effect due to surface charge is negligible to avoid the difficulty in estimation of the electric field effect. In this microregion, there is ion interaction between both the surface and the electric potential effects (which are induced by the surface charge).


The Surface Charge

The colloidal behavior depends on how the surface

charge originated. There are two type of colloidal charges. The first type has the charges due to crystalline imperfection, such as isomorphic substitution of Si4+ by Al3+ or by 2+
other cations such as (Mg ) having a lower charge. This type of charge is found in the crystal lattice of clays such as montmorillonite and vermiculite. In these cases, the charge density is constant per unit of surface area. The second type occurs where the surface charge may be created by preferential adsorption of a certain ion, such as that by hydroxyl or phosphate ions. In this case, the charge arises at the exterior edge of crystals or lattice, thereby inducing a constant surface potential (van Olphen 1977; Bolt 1976).









Gouy-Chapman and Stern Theories of the Double-Layer

From Gouy-Chapman theory, the charge density can be deduced considering the electroneutrality condition (van Olphen 1977; Sennet and Olivier 1965):


a = -f pdx (60)
-akt
= -. exp(zieo/2kT) - exp(-zieTo/2kT)] (61) where p is the space charge density (or net sum of positive and negative ion concentration, a is the surface charge, and is the surface potential. The Gouy-Chapman theory gives an over-estimation of the double layer capacity, namely

K = (8z2 e2a/E kT) 1/2 (62) where K is the reciprocal of the double layer thickness, a is the ion concentration, e is the electron charge, z is the valence of the ion, and c is the dielectric constant.

Stern recognized the importance of the ion sizes near the surface. He proposed that the counter ions could be divided between a diffuse layer and an immobile surface layer of thickness 6 able to contain a maximum number of counter ions per unit of surface area. This may be expressed as
no m +p(T-- exp(zeT6/2kT)-exp(-zeT6) (63) wen+Aexp(-ze6/kT+he cs where a mis the charge corresponding to a monolayer of








counterions, q is the van der Wall energy, A is the frequency factor, T6 is the electric potential at the Stern Layer.


Zero Point of Charge (ZPC)

On a metal oxide surface, charges are created by the adsorption and desorption of H+ or OH ions which are affected by their concentrations in solution. The Parks and de Bruyn (1962) model is

OH + OH OH

+oH1- +OH - + ( 4 M A = M C (64)

\OH2 +H+ \OH +H+ OH


where M is Al or Fe with A- and C are the associate anion and cation, respectively. From the Eq. (64), it can be seen that there is a pH at which the net surface charge is zero; this point is termed the zero point of charge (ZPC).

For metal oxides, such as described above, H+ and OH ions are potential determining ions. By approximation to the Nernst equation, Keng and Uehara (1974) reported that

PT H+
SR T in H-- (65)
0

where R is the gas constant, T is the absolute temperature, F is the Faraday constant, and 4o <<25 mV; the Gouy-Chapman double layer equation is reduced according to









o =2- = 2- (0.059)(pH - pH) (66)


Thermodynamics of Adsorption Effect of the Electric Field

In this system introduced in Fig. (1), the compartments A, B, and C can be considered as being three phases of a particular ion at different potential energy levels. In each phase, molecules are at different energy levels and varying within a range of potential characteristics of each phase.

Around a colloid center, the ions exhibit a Boltzman type distribution:

c = b exp [-F(T c - T b)/RT] (67a) b = a exp [-F(b - T1a )/RT] (67b)


Supposing that Ta = 0, it can be shown that

T, Tc - RT n -c (67c)
0 c F a

where To is the surface potential, Tb and T1a are electric potentials that is characteristic of the phases B and A, respectively.


Free Energy as a Function of Distance

From the outer limit of the Stern layer, both Gaussian and Gouy-Chapman concepts can be applied to a particular ion. The movement of an ion (or molecule) results from a succession of collisions that may move it at random in a











positive or negative direction. These ionic concentrations follow the Gaussian distribution:

a M exp(-x2 /4Dt) (68)
2(flDt)

where M is the amount of substance deposited at the plane when x = 0 at the time t = 0, and D is a constant. Assuming that the Gouy-Chapman distribution holds in the following equation:

T= T exp(-Kx) (69) where the free energy can be expressed as

G. =G0 + zFT + RTln a (70a)
1,

and
2
G. = G� + zFT exp(-Kx) + RTln N RTX (70b)
im 0 2(Dt)i/2 4Dt

At constant time t, this expression becomes

d ZG F Ke-Kx RT (71)
dx 0 Z~ - TD (1

The limit between the physically adsorbed layer and the completely free ion in the phase A may be defined when dG/dx=0. If we set y = -Kx and y = -RT/2Dt/zFT oK, then the expression can be written

y
e y 0 (72) and become

y + 2y(l - .) + 1 = 0


where









(ey = 1 + y + 1/2 y2 + . . . )

When this quadratic equation gives two roots, they may be indicators of the transition states between different degrees of adsorption.


Free Energy for Irreversible Fixation

The active transport of an ionic substance against the gradient potential is determined by the difference between the total potential within each phase (A, B, and C). The maximum energy, other than expansion work for a change of ion activities, is expressed by G. at constant temperature i ,m
and pressure. The transfer of ions between the phase B and C is due to a potential energy minimum for phase C written
c B
as Gc and maximum for phase B written as G. . In this
'm 1,m case

GC z.FTc + X. (73)
1,m 1 1

and, correspondingly,

G. = z.FB + RTln b (74)
1,m 1

where G. is also the Gibbs free energy of the ion i per
1 ,m
unit mole, Xi is defined as the minimum energy for irreversible fixation (including only the electro-chemical energy), F is the Faraday constant, 11' and T are the electric potential in the phases B and C, and the term RTln b is the chemical potential. The transfer of ions between the phases B and C is governed by the difference in total energy between the two phases,









AG. -46az.F + RTln s - + RTw2(t - t ) - X. (75)
IM1 k 0 i2o

where T= T C - T B = 0- T 6 = 4H6a which is the Gaussian equation for a molecular condensor. This by substitution gives
AG. 4fl6oz.F + RTln - -2 X. (76)

1lm 1 2 1

When the conditions are at equilibrium, or AG = 0 and t = to, then

X. = - -46az.F + FTln s k2 (77)
1 1 k2

It is assumed that the minimum energy Xi is independent of the state of equilibrium, then for a given colloid surface which has a particular surface charge, the free energy change AG. is reduced to
1lm
AG RTw (t - t ) (78)
1,m '2 o


Relationship of Surface Charge to Surface Potential

During the overall transfer of ions from the phase A

to phase C at equilibrium conditions, after substituting the Xi value of Eq. (77) and assuming 'A = 0, the relationship becomes

411
do d C (79)


where 6 is the Stern layer thickness, a is the surface charge, F is the Faraday constant, and a. is the ion valence.
i









Surface Tension and Specific Adsorption

In the interfacial colloid-liquid region, some ions

may enter into the coordination shell of surface atoms. As a result, there is a modification of the colloid surface (A) (expansion or, contraction). This work required to expand or contract the surface when divided by the change in surface area is the surface tension (y) (Atkins 1978). The following series demonstrates the relationship:

GC = z.FIc + Ay (80a)
1 1

G = z.Fc + X.c (80b)
1 1 1

Differentiating the Eq. (80a) and Eq. (80b) gives

dG.C z.FPdc + z.FcdT + ydA + Ady (81a)
1 1 1

dG = z.Fdc + z.FcdT + X.dc + cdX. (81b)
1 1 1 1 1

However, it is known that

dG = z.Fdc + X.dc (82a)
1 1 1

dG 1 = z.Fl'dc + ydA (82b)
1 1

By comparing Eq. (81a) to Eq. (82b) and Eq. (81b) to Eq. (82a) it follows that

cdX. = Ady (83)
1
The differentiation of Eq. (77) gives dX. -46flz.F.do, so it can be deduced that
-4flHtz.
dy 1 (84)
J5 A c













CHAPTER IV
MATERIALS AND METHODS


Goethite Preparation

Three goethite preparations were made by mixing separately an equal volume (50 ml) of 1 N Fe(Cl)3 with a similar volume of 2 N NaOH, 1 N NaOH, or 0.5 N NaOH in order to have OH/Fe mole ratios of 6, 3, and 1.5, respectively, in the mixtures. The suspensions were allowed to age for one week to induce crystal formation and growth. The addition of phosphorus to the goethite preparation was made for P/Fe = 3.2 ratio both at the beginning and by adjustment at the end of ageing process (one day of phosphorus reaction is allowed). In order to examine the effects of phosphorus concentrations on goethite crystallization, different levels of phosphorus in 2 N NaOH were mixed with 1 N Fe(Cl)3 and allowed to age. After one week of ageing the suspension at either room temperature or at 550C, the samples were washed by dialysis against distilled water during one week; the water was changed after intervals not greater than 12 hours. At the end of the ageing process and washing, the samples were freeze-dried.

The existence of hydroxyl deformational vibrations in
-1
the region 1200 to 1000 cm was investigated by two methods:









1) goethite and phosphated goethite were digested in D 20 for 24 hours to replace the surface OH by OD, or 2) phosphated goethite samples were equilibrated with different salts (0.1 N KCl, 0.1 N KNO3, and 0.1 N Na2SO4) and water to displace surface phosphates. The samples were separated from solution and dried for the infrared spectroscopy studies.

An attempt was made to prepare FeOOD by digesting three times 0.9 g of FeCle .6H 2) in D 20 and drying the suspension in order to make FeCl 3. 6D2 0. The residues were dissolved in 10 ml of D 20 and mixed with 10 ml 2 N NaOD; the suspensions were then aged and dried as described above. Infrared and X-Ray Diffraction Techniques

One milligram of the freeze-dried sample was mixed with 400 mg KBr to make pellets samples used for the infrared study. The infrared spectra are obtained by using PerkinElmer 567 grading infrared spectrophotometer.

Prior to X-ray diffraction analysis using a general electric XRD700 instrument, and thin films of goethite samples were made on glass slides and allowed to air dry. Some Factors Affecting the Kinetics of Adsorption-Desorption

The adsorption or desorption isotherm is obtained by plotting the amount of ion adsorbed or desorbed per unit weight of adsorbent versus the solution concentration. The









solid and the solution were equilibrated through agitation for selected periods of time and temperatures. The equilibrium solution was removed after centrifugation and the phosphorus concentration determined through blue color development as ascorbic acid molybdophosphoric complex.


Effects of pH

The adsorption for different periods of times (0 to 72 hours) was done by equilibrating goethite or soils with phosphorus solution (Sg P/ml) or (20 pg P/ml) adjusted to pH 2, 4, 8, 9.5, and 11. Solid phase was separated by centrifuging at 5000 revolutions per minute (rpm). Phosphate desorption was carried out on samples which reached the equilibrium state. The desorption was conducted in water, adjusted to the above pH, during 20 minutes, 6 and 12 hours. The pH of the solution was measured by placing the glass electrode in the supernatant after centrifugation. Effects of the Supporting Electrolytes

The effects of the type of cations on the adsorption

were studied by mixing 20 mg of goethite with 20 ml each of the phosphate solutions containing 1, 3, 5, and 10 vg P/ml in 0.01 M salts of NaCl, CaCl2, and AlC13, each in separate experiments. With soils, the phosphate solution concentrations used were 0, 10, 20, 40, and 80 pg P/ml. The extent that each of the cations (Ca, Na, and Al) contributed to the








degree of reversibility of the adsorbed phosphate was studied by conducting the desorption with solution of the corresponding electrolyte (as in the adsorption) but without phosphate. A comparison of the phosphate desorption by different anions in various salts (KCl, K2So4, KN03, KCI04, and NaOH) was made for goethite-P solution (one gram of solid with one liter of salt solution where the goethite had previously been treated with 5 vg P/ml). Time Effects on Adsorption and Desorption

Adsorption of phosphates on goethite-solution containing 0, 1, 3, 5, 7, and 10 jg P/ml, and on soil-solution containing 0, 5, 10, 20, 40, and 80 vg P/ml was investigated after equilibration during different periods of time ranging from 1/10 hour to 80 hours. At the end of each reaction time, the solid-solution was centrifuged and the phosphate concentration remaining in the supernatant solution determined. The solid/solution ratio was 1/1000 for goethite-solution and 1/20 for soil-solution. Kinetic studies of phosphate adsorption were done at different conditions: 1) when different types of electrolytes were used, 2) for one electrolyte (CaCl2) at different concentrations (.01 N Ca,

0.1 N Ca, 1.0 N Ca), and 3) at pH 2 and pH 10. After the adsorption proceeded until the equilibrium state was reached, desorption was conducted during 20 minutes, 6 and 12 hours in order to determine the minimum time required for complete









desorption of this form of phosphate. This minimum time was also employed during desorption of phosphate when using either a desorbing solution at selected pH (2 to 11) or for different electrolytes. For those soils where it was observed that phosphate sorption did not follow the Langmuir isotherm, an attempt was made to determine the extent of changes in aluminum and iron phosphate by using the method of Peterson and Corey (1966). For aluminum phosphate, one gram of soil sample was washed with 2 N NaCl to remove exchangeable cations, the suspension was centrifuged and the supernatant solution discarded. Then 20 ml of 0.5 N NH4 F at pH8.2 were added and the suspension was shaken for an hour and then centrifuged for the P determination. For the iron phosphate fraction, the samples were washed with 2 N NaOH solution, centrifuged, and the supernatant retained for P determination.


Surface Charge as Affected by Phosphate Adsorption

The method of Lavardiere and Weaver (1977) was used to determine the net electric charge. The procedure was to add 20 ml each of 0.01 N CaCl2 and 1.0 N CaCl2 solutions per one gram of soil sample or 20 mg of goethite. Subsequent titrations of the suspensions were made with 0.01 TI HCl or 0.01 N NaOH, by adding 0.1 to 0.3 ml at a time from a microburet at 2-minute intervals. Continuous stirring was maintained and the pH read before the addition of either base or acid.









A blank titration was made for the same volume of CaCl2 solution. The amount of H + or OH- adsorbed at a given pH was calculated by the difference between the amount of H+ or OH- added and that required to bring the blank solution of the same volume and salt concentration to the same pH as the soil or goethite.


Soil pH, Iron Oxide, and Extractable P

The soils used were obtained from Georgia (Cecil, Ap

horizon), Colorado (roadside cut near Ft. Collins), and Kenya (latosol, sampled at 15-30 cm, near Kabete).

The pH was measured in 1:1 soilto solution suspension

for 10 g of soil with 10 ml of H20 or 10 ml of 1 N HCI, using a glass electrode. Iron oxide was determined by the dithionite-bicarbonate extraction and colorimetric determination of Fe as the ferrous orthophenonthroline complex. Extractable P was determined by three of the methods outlined by Ballard (1979). These data are shown in Table 1.









Table 1. Soil pH, iron oxide and extractable P.


Soils Georgia Colorado Kenya pH(H20) 5.88 8.17 6.20 pH(KC1) 4.95 7.44 5.75 % Fe203 2.24 0.67 3.95 P, ppm 0.05 N HCI + 0.025 N H2 S04"

5.60 0.30 0.40 P, ppm 0.03 N NH F + 0.1_N HCI* 13.00 7.40 1.40 P, ppm 0.5 M NaHCO3

2.80 1.20 1.60


*Reagents described by Ballard (1974)













CHAPTER V
RESULTS AND DISCUSSION

Goethite and Phosphated Goethite Studies by InfraRed Spectroscopy

After the mixing of FeCl3 and NaOH solutions, some

precipitates formed, but the particles formed at precipitation were not yet crystalline. The changes in precipitates to crystals originated from a discrete center or crystal nuclei (twinned and acicular crystals) produced by different mechanisms (Atkinson et al. 1968). The conditions governing the formation and nature of these crystal nuclei are not well known. Apparently, the number of crystal nuclei was increased by both temperature and hydroxyl ion concentrations. Iron and hydroxyl ions would be attracted to the centers of crystal growth as they lose their energies. As iron and hydroxyl ions were involved in the formation of the crystal, ageing reactions follow the sequence of lepidocrocite to goethite (Murphy et al. 1975). Another way would be that ferrhydrite forms first or just goethite was formed. The growth of the crystal occurred as the ions change phase during deposition from the solution. 1su (1972) found that the removal of hydroxyl ions from the solution resulted in a drop in pH value of the solution as the time increased.









Since the ions are changing phase from the solution (higher degree of randomness) to the solid phase (lower degree of randomness), the entropy change during the process from solution to crystalline phase must be negative. The entropy of ions within the crystal would be lower than those at the crystal surface. From infrared spectra, Kiselev and Lygin (1975) believed it was possible to estimate from adsorption entropy what was the degree of freedom of the adsorbed molecules.


Goethite Structure Identification by Infrared


The 4000 - 2000 cm-1 Region

The goethite structure was interpreted through the

identification of OH and FeO vibrations. The hydroxyl ion vibrations occurred at two strong bands in the 3700 - 2000 cm- region, centered at 3400 cm- and 3200 cm-1 (see Fig. 2 and Fig. 3). According to Nakamoto (1978), lattice water absorbed at 3550 - 3200 cm- , he reported that the strong band centered at 3400 cm-1 was due to antisymmetric and symmetric OH-stretching of water. It could easily be recognized that the 3400 cm-1 band was not characteristic of goethite crystal structure since this band appears even when no goethite is present, as determined by X-ray diffraction which showed no crystalline material. It is of interest to observe that the intensity of the 3200 cm-1






50




OH/Fe 6



A


A'











B

H- B'





C')














3500 3000 2500

Frequency (cm-') Fig. 2. Infrared bands of (A) goethite; (B) phosphate added
to goethite at the end of ageing; and (C) phosphate
added to goethite at the beginning of ageing.
Curves A, B and C are for 220 C and curves A', B',
and C' are for 550 C.





51




A OH/Fe 6






A''








B'








C-)


C<












1200 1000 800 600 400


Fig. 3. Infrared bands of (A) goethite; (B) phosphate added
to goethite at end of ageing; (C) phosphate added
to goethite at beginning ageing. Curves A, B, and C are for 220 C, and curves A', B', and C' are for
550C.









band decreased as the OH/Fe ratio in the ageing solution decreased. The band centered at 3200 cm- appeared only if goethite structure was present. As a result, it was con-i
cluded that the 3200 cm band corresponded to structural Fe - OH stretching.
-i
The 2000 - 3000 cm Region

After goethite was freeze-dried, not all the adsorbed water was removed. Nakamoto (1978) indicated that the relative velocity of the oxygen nucleus compared to that of hydrogen nucleus is small. This meant that the surface binding of water through the oxygen atom to goethite would not induce a significant change in the overall water vibration which could be observed at around 1620 cm-1 as depicted for water structure below



1' 4
0 0 0
H H H H H H Case v1 Case 2 Case v3


The above three normal modes of vibration in H20 are infrared active. The bending vibration v2 is centered at 1620 cm and the water stretching bands (v1 and v 3) vibrate in the 3400 cm region.








-i
The 1300 - 700 cm Region
-1
The two main bands at 890 and 790 cm would be assigned to Fe-OH bending vibration of the structural hydroxyls. Busca et al. (1978) supported this finding by observing that the structural OH in c-FeOOH disappeared upon heating to
-I
form a-Fe 20 3* The appearance of the 890 and 790 cm bands always indicated goethite crystallization, which can be observed in Fig. 3, 4, 5, 6, 7, 9, 10 and 11. From the above mentioned figures, it was observed that for each OH/Fe ratio, the way in which phosphate was added had an effect on the degree of goethite crystallization. From Fig. 5, as reported by Farmer and Palmieri (1975), typical goethite could be identified by the Fe - OH bending vibration at 890 cm and 790 cm-. In Fig. 3, where OH/Fe = 6, there are also strong bands at 890 and 790 cm- in all cases, except when phosphate is added at the beginning of ageing of the suspension at 550C. In the latter case, the increase of temperature to 551C favored preferential phosphate binding to the iron which was observed by the strong vibration at 1000 cm- . As shown in Fig. 3c, phosphate was within the goethite structure because there was both the P - O(Fe) vibration at 1000 cm- and the surface binuclear (FeO),POOH
L
vibrations at 1190, 1100, and 1030 cm-. When OH/Fe 3, it was not possible to have phosphate within the crystal because there was no crystallization at room temperature. However































< B'




C)C














3500 000 2500 Frequency (cm-I)
Fig. 4. Infrared bands of (A) goethite; (B) phosphate added
to goethite at end ageing; and (C) phosphate added to goethite at beginning ageing. Curves A, B, and
C are suspensions at 221C, and curves A', B', and
C' are for suspensions at 550C.










I I I 1 I


Goethite


Lepidocrocite


I I I I I


1600 1400 1000


800 700


Wave Number (cm')




Fig. 5. After Farmer and Palmieri (1975). Infrared bands of
Goethite and lepidocrocite.


6 0
600 500


I I I


I I



































C-)


z <
















1200 1000 800 600 400) Wave Number (cm-I)
Fig. 6. Infrared bands of (A) goethite; (B) phosphate added
to goethite at end ageing; and (C) phosphate added
to goethite at beginning ageing. Curves A, B and
C are for suspensions at 220C and curves A', B',
and C' are for 550C.

































(9 H



C'












1200 800 600 400 Wave Number (cm-I)
Fig. 7. Infrared bands of (A) Fe hydroxide material; (B)
phosphate added to Fe hydroxide material at end of
ageing; and (C) phosphate added-to Fe hydroxide
material at beginning ageing. Curves A, B, C are
for suspensions at 220 C and curves A', B', C'
are for suspensions at 550 C.








there was weak crystallization at 550C that favored some phosphate surface binding (see Fig. 6b). In OH/Fe = 1.5 suspension, after one week of ageing either at room temperature or at 550C, no goethite formation was observed. There was a single, strong vibration at 1030 cm-1 when phosphate was added at the end of the ageing, but the vibration was at 1000 cm1 when the phosphate was added at the beginning of ageing. The 1030 cm- vibration, which is the P - OH bending vibration, indicated the excistence of phosphate at the
-1
surface, and the presence of the 1000 cm vibration from the P - O(Fe) showed phosphate directly bound to iron when phosphate was added just prior to the ageing process (Parfitt et al. 1975).


Goethite Identification by X-Ray

The X-ray diffraction peaks for goethite were at 4.19 2.70 R, and 2.45 R. In all cases, the 4.19 R is the most intense peak while those at the 2.70 R and 2.45 R were weak, Fig. 8. The presence of the 2.70 R spacing indicated that some hematite might be present. However, because the 2.45 peak was present, it was believed that considerable amounts of goethite existed (Schwertman and Taylor 1977). According to their studies, it would be possible for hematite to exist along with goethite because the hydration of hematite would yield goethite with a standard free energy (AG0) of the reaction varying from -0.2 to 0.4 kcal/mole.








Angstrom Spacing ( )


35 30 25 21


Degree (20)
Fig. 8. X-ray diffraction patterns of (A) goethite, (B) phosphated
end of ageing, and (C) phosphated goethite at beginning of


goethite at ageing.


# I









Factors Affecting Goethite Crystallization Effect of OH/Fe Ratio

The above results indicated that the bands at 3400 cm
1 -i -i -i , 3200 cm , 890 cm , and 790 cm were characteristic of the appearance of goethite. For the same ageing period, an increase in OH/Fe favored goethite crystallization. When the OH/Fe is 6, the bands at 3200 cm-1, 890 cm-1, and 790 cmwere strong, indicating that goethite structure was wellformed. Where OH/Fe is 3, after a week of ageing, goethite structure was apparently present only if the suspension was kept at 551C, even then the degree of crystallization was less intense at OH/Fe = 3 than it was where OH/Fe = 6. This weak crystallization was suggested by the weak band at 3200 cm-1 (which took the form of a shoulder) and by the
-I -i
less pronounced intensities of the 890 cm and 790 cm bands. However, when OH/Fe is 1.5, there was no goethite crystallization even when the suspension was aged at 551C. The above conclusion disagreed with that of Atkinson et al. (1974) who claimed made goethite was in a suspension at 280C where OH/Fe is 2.0 when aged for 50 hours. Phosphated goethite

The application of phosphorus weakened the goethite

structure because some phosphate was probably incorporated within the lattice. In Fig. 4 and Fig. 6 where OH/Fe is 3,








the addition of phosphate to goethite suspension resulted in
-I-I -i
bands at 3200 cm , 890 cm and 790 cm which were weaker than those for OH/Fe = 6 after ageing. When OH/Fe = 6 in the suspension, hydroxyl ion concentration was high enough to form the necessary bonds for the goethite structure to appear. The presence of phosphate on the goethite surface could be recognized by the appearance of bands in the 1200
-l
1000 cm region, but vibrational hydroxyl deformations also could occur in this same region. Parfitt (1979) stated that the P = 0 bond had stretching vibrations in the 1190 cm1 and
-I
the 1030 cm region.

When the phosphate was added at the end of the goethite ageing, there was an appearance of P = 0 vibration at 1190 cm and the 1030 cm due to P - OH vibration in agreement with work by Parfitt (1979). At high temperature (550) even when OH/Fe = 6, the phosphate is preferentially bonded to the iron, thereby reducing the capacity of hydroxyls ions to bind freely with iron to form goethite. As a result, there was direct binding of phosphate to iron which was confirmed by the strong band at 1000 cm- assigned to P - O(Fe) and by the lack of the Fe - OH bending vibrations at 890 cm and 790 cm1. This is confirmed by spectra given in Fig. 3. In Fig 9, Fig 10 and Fig 11 where suspension the OH/Fe ratio was either 6, 3, or 1.5, the increase in phosphate concentrations relative to the iron (P/Fe = 0.032,





62





OH/Fe = 6





P/Fe 0. 32











C.)

'-4



















120 1IC0 )QO0 600 4O0





Fig. 9. Infrared bands of phosphated goethite at beginning
of ageing for suspensions of various P/Fe values at
OH/Fe = 6.





r3










OH/Fe 3. 0



f





P/Fe =0. 03










C
z z

















1200 1000 800 600 400 Wave Number (cm-) Fig. 10. Infrared bands of phosphated goethite at beginning
of ageing for various P/Fe ratios at OH/Fe = 3.0.







































L










1200 1000 800 600 400 Wave Number (cm-1)

Fig. 11. Infrared bands of phosphated goethite at beginning
of ageing for various P/Fe ratios at OH/Fe = 1.5.








0.32, 3.2) induced an increase in the intensity of P - O(Fe) and P - OH stretching vibration at 1000 cm- and 1030 cm-I, respectively. When P/Fe was greater or equal to 0.3, these two bands overlapped, showing only one very strong band at 1000 cm1 .

In order to identify the vibration of hydroxyl deformation assumed to be in the 1200 - 1000 cm region, two methods were used: 1) goethite and phosphated goethite were digested in D20 for 24 hours and 2) surface phosphate was desorbed by different anions and water. After the samples were dried, the infrared spectra were similar to those shown in Fig. 12 and Fig. 13. The weak bands due to OH deformation were displaced by OD and only the P - O(Fe) stretching vibration at 1000 cm-1 and those for the Fe - OH bending at 890 cm- and 790 cm persisted. Evidently, it could be assumed that the bands at 890 cm- and 790 cmwere due to structural Fe - OH bending which cannot be affected by digestion in D20 as long as the initial goethite maintained its structure. In a further study, the evidence for hydroxyl deformations was examined when the phosphated goethites were desorbed by different salts (KCl, KNO3, Na2SO4) and water. After either 0.1 N KCl solution or water desorption, hydroxyl deformations were not removed in the 1200 cm - 1000 cm region. With 0.1 Na2S04 solution used for desorption, the infrared spectra showed strong




66



















(2)


UI






















1200 1000 800

Wave Number (cm-1



Fig. 12. Infrared hands of 1) goethite digested in 1) 2 and
2) phosphated goethite digested in D 20.2





67








(1)






()












z

F
















1400 1200 i000 C0

;';ave Number (cm-1) Fig. 13. Infrared bands of phosphated goethite after desorption by (1) 0.1 1 KC1, (2) H20, (3) 0.1 _ KN03,
(4) 0.1 N Na 2SO4








bonds of SO2- ions in the 1200 - 1000 cm- region. When
0.1 N KNO was used for desorption, NO3 appeared to
3 displace some surface phosphates so that the hydroxyl vibration in that region was reduced (Fig. 13). It can be said that NO3 and SO ions have the ability to desorp readily displaceable phosphate at the surface. However, sulfate ions could not be used to provide evidence for this type of phosphate desorption by infrared since sulfate ions have strong vibration in the same region as that of surface-bound phosphate. Parfitt et al. (1975) assigned the appearance of weak bands in the 1200 - 1000 cm region for non-phosphated goethite to the Fe - OH deformational vibration which can partially overlap the region of P = 0 stretching and P - OH bending vibration.

The question arose about the type of hydroxyl groups

that are displaced when phosphates are added to the goethite suspension. Russel et al. (1974) indicated that there were three types of hydroxyl groups: 1) where OH is singly coordinated to an Fe atom with hydrogen bond interaction with another atom, 2) where OH is coordinated to two Fe atoms and 3) when OH is singly coordinated to an' Fe atom. Parfitt et al. (1975) considered that only type 3 of OH was displaced by phosphate while OH of the type 1 and 2 are unreactive. The types 1 and 3 of OH, they considered as the structural OH with three vibrations at 3200, 890, and
-I
790 cm.









OH OH OH Fe Fe Fe Fe

I I I I
OH OH OH OH Fe Fe Fe I I 1 [001] face type 3 ofOH -* OH OH OH Fig. 14. The 001 face of goethite lattice. (After Bragg
and Claringbul 1965).


The formation of binuclear bridging resulted from the displacement of two adjacent type 3 hydroxyls



OH OH OH Fe Fe Fe Fe OH OH OH OH Fe Fe Fe
I I
O 0 OH

P

0 0


Fig. 15.


The 001 face of phosphated goethite. Phosphate Adsorption and Desorption Studies


Effects of the Initial Concentration

The phosphate adsorption isotherm is obtained by plotting the amount of phosphorus adsorbed (vg P/g of adsorbent)









against the equilibrium concentration of phosphorus in solution. The shape of each isotherm curve is altered by the relative amount of phosphorus adsorbed at each equilibrium concentration and reaction time. For a particular time of adsorption reaction, the amount of phosphorus adsorbed increases with the initial concentration but the percentage of phosphate sorbed decreases.

For the goethite-P solution system, the shapes of the phosphate adsorption isotherm curves were affected by the relative amount of phosphorus adsorbed at each equilibrium concentration (Fig. 16). Each curve was composed of three parts: one at low solute concentration (< 0.25 pg P/ml), a second part at higher solute concentration (0.25 to 1.5 vg P/ml) where the isotherm becomes convex, and a third linear part at higher solute concentration (> 1.5 og P/ml).

However, in the soil-solution system, phosphate adsorptive capacity depended on soil characteristics which affected the shape of the isotherm curves. The Kenya soil and Georgia soil phosphorus sorption curves apparently had two portions, the first part which is a curved portion (< 5 vg P/ml) and a second or linear portion (> 5 pg P/ml). In Fig. 17, the apparent lack of an initial linear portion for the curve found for Kenya soil suggests need for more adsorption data at low equilibrium concentration (< 3 g P/ml). For the Georgia soil, (Fig. 18), the adsorption maximum is low compared to































3 hours


0 3 4 5 6 Equilibrium Concentration (iig P/ml) Fig. 16. Adsorption isotherms as affected by the reaction times for goethite-solution

systems.


-' 3 4-J
4
0)
0



S2




0
-4


0 Cl)










0













































U 10 20 30 40 50 Equilibrium Concentration (jig P/ml)











Fig. 17. Adsorption isotherm for the Kenya soil after 24
hours of reaction time.
























0
4



0
0









0 10 20 30 40 50 60 70 E .r Eqiiru Cnetaio w /4


Fig. 18. Adsorption isotherm for the Georgia soil after a reaction time of 24 hours.



























4






- 2






0
0 10 20 30 40 50 60 Equilibrium Concentration (vg P/mi) Fig. 19. Phosphate adsorption isotherm for the Colorado soil after 24 hours
of reaction time.









that for the Kenya soil, which means that for the same initial concentration in solution more phosphate ions are present in the Georgia soil than in the Kenya soil. For the Colorado soil, (Fig. 19), there is a slight change in slope for the initial and final linear party of the absorption isotherm. For the initial range of equilibrium concentration ranging from 0 to 25 pg P/ml, the amount adsorbed increased linearly as the equilibrium concentration increases. This relationship was 20 jig P adsorped per gram of soil for each pg P/ml. Muljadi et al. (1966) noted that their isotherm curves were also linear initially with a curved transition to the second linear portion. They ascribed the initial and curved portions to phosphate exchange with OH of Al(OH) located on the clay edge surfaces. They do not give a clear explanation of the mechanism of adsorption reaction for the third part of the curve but postulated that the final linearity of the isotherm indicated that the number of adsorption sites remained constant even though the amount of phosphate adsorbed increased.

The time required for the reaction to reach equilibrium decreased as the initial concentration decreased, (Fig. 20). If the initial concentration was less than or equal to

5 pg P/ml, the steady state of reaction was reached in less than 10 hours. Twenty hours of reaction were required to approach the equilibrium state with an initial concentration

























-,4
-j 5 P/mi


4




C 0


. 2


0



0








0 10 20 30 40 50 60 Reaction Time (hours) Fig. 20. Effects of the initial P concentration on the phosphate adsorption by

goethite-solution (1 g/lO00 ml).









greater than or equal to 5 vg P/ml. The above observations showed that low amounts of phosphorus were almost instantaneously adsorbed onto the synthetic goethite. Because the initial phosphate potential on the solid phase was low or nil relative to the phosphate potential in solution, the phosphate flux from the solution was therefore high. This movement of phosphate to the solid surface continued for a longer period of time if the initial concentration (or phosphate potential) in solution was high.


Effects of the Supporting Electrolyte Type of Electrolytes

The adsorption isotherm of phosphate on geothite was examined using molar concentration of the electrolytes, NaCI, CaCl2, and AlCI3, as given in Fig. 21. Where the initial concentration was greater than 0.5 Vg P/ml, the salts (electrolyte) gave a significant effect on the amount of phosphorus adsorbed. The adsorption at each concentration decreased in the order iM AlCl3>lM CaCl2>lM NaCl. It was noted that the adsorption of phosphate using water as supporting medium gave the same adsorption isotherm as that using 1 M NaCI solution. For equilibrium concentration greater than 1 ug F/ml, the amount of P sorbed in 1 N AICl3 and 1 M CaCl2 was greater than that sorbed in 1 M NaCI by a factor of 1.6 and 1.5, respectively.

















5

0


-4

C
U3 3





0
o 2











1 2 3 4 5 Equilibrium Concentration (Dg P/mi) Fig. 21. Effects of the type of supporting electrolyte on P adsorption on
goethite-suspension (1 g/1000 ml).









Table 2. Type of salt effects on the P adsorption on
goethite at 5 ig P/ml.


Electrolyte P adsorption (pg P/g)

1 M NaCl 3780 1 M CaCl2 5600 1 M AlCl3 5900




Ryden and Syers (1975) reported that, where some soils have a final concentration above 0.1 vg P/ml, the P sorption in 10-2 M Ca solution was 1.5 to 2.5 greater than the sorption from water. From the isotherms for P sorption by goethite, the data were arranged as the plots of a/b versus a, where a is the equilibrium concentration, b is the amount of phosphorus adsorbed per unit weight of adsorbent. The linearity of the plots (Fig. 24) confirmed the Langmuir type of adsorbent. When using Eq. (14c) the adsorption maximum

(s) and the adsorption energy constant (k) are calculated as shown in Table 3.
















Regression equations
a/b = 7.47:'.i0-4c a /b = 4 .39*- -,10 4 a/b = 2.60110-4 a/b = 2.24*10-i


+ 1.13:':10-3 + 5.21*104 + 1. 0910f 5 + 4.84":10-


0 1 2 3 4 5 6 Equilibrium Concentration (pg P/ml)









Fig. 24. The transformed Langmuir equations for phosphate
sorption by goethite as affected by reaction times.
a is equilibrium concentration and b amount adsorbed.









Table 3. Effects of three electrolyte salts on the phosphate sorption maximum and sorption energy constant for goethite.


Electrolyte Adsorption Maximum Sorption Energy Constant (Pg P/g) (ml/pg P)

1 M NaCI 4300 3.8 1 M CaCl2 6900 2.3 1 N AlCl3 12700 1.1



The increase of phosphate sorption maximum, (Table 3), was accompanied by a decrease in the apparent sorption energy constant. The increase in phosphate sorption may be due the increase in the cation charge of the electrolyte favored adsorption. A probable way in which the cation charge enters into the phosphate adsorption reaction was through cation bridging:



Fe 0
n+I
S---0 - P 0 (85)
OH
Fe 0

so that where M is Na, then n = 0, or if N is Ca then n = 1, and if M is Al then n = 2. As the charge on the cation increased, there was a greater attraction between the cation









on the goethite surface and the phosphate ion. The work required to bring the phosphate ion to the cation evidently decreased as the valence increased. The increase in metal valence favored a higher probability that the phosphate was maintained in the goethite-P solution interface, so that the energy for adsorption decreases.

Both aluminum and calcium ions also probably reacted with phosphate in solution to form new phases (precipitation), inducing thereby a decrease of phosphate in solution. Some of the compounds which might precipitate in solution depended both on phosphate concentration (molarity) and pH. These systems can be written as follows:

For CaHPO4, the pH and H2PO4 relation is

PH2PO4 = pH - 3.14 (86)

For Al(OH)2H2 P04, the pH and H 2PO relation is

PH2PO 4 = pH + 10.7 (87)

For Fe(OH) 2 H2PO4, the pH and H 2PO relation is

PH2PO4 = pH - 10.9 (88) The plots of PH2PO 4 versus pH gave the solubility diagrams of the compounds as illustrated by Lindsay and Moreno (1960). At any pcint (pH, PH2FO4) above the line for a selected compound, precipitation is expected while below the lines dissolution of the corresponding solid phase will occur. In our study, the pH ranged from 6 to 7 so that with








a phosphate concentration of 5 iig P/ml or 1.72xl0-4 moles of H2PO4 per 1 liter, we would have PH2PO 4 = -log H2P04 = 3.8; hence, no important amount of precipitation of the above solid phase was expected. Other complex compounds involving combination of Ca and Fe phosphates or Al and Fe phosphates could precipitate near or on colloid surfaces of the goethite or the soils.


Effects of Electrolyte Concentration

Salt concentration had a marked effect on the amount of phosphate sorbed during reactions at various periods. The time at which equilibrium was reached was evidently not affected by increasing the salt concentration (Fig. 22). When the initial concentration is 5 vg P/ml, at equilibrium state the phosphate sorption in 1 N CaCl2 and 0.01 N CaCl2 is 1.3 and 1.2, respectively, greater than that found for water system without salt. Van Olphen (1977) pointed out that increasing the electrolyte concentration not only caused compression of the diffuse part of the double layer but also some ions, as counter ions, shift from the diffuse layer to the Stern layer. As a result, the ion concentration of the diffuse layer decreased and more adsorption took place. This concept was supported by Ryden and Syers (1975) who found that equilibrium phosphate concentration remained the same in both 10-2 M Ca and 10-I M Na systems. Donnan theory could be used to explain the above phenomenon. The inner solution possibly ranged from the solid surface to the








































0 10 20 30 40 50 Reaction Time (hours)











Fig. 22. Effect of the supporting electrolyte concentration
on the kinetics of P adsorption by the goethitesolution system (1 g/1000 ml).









Table 4. Phosphate desorption from goethite by different
anions.


Anions (0.1N)

C]- so 2 NO - ClO0*1 OH - HO0 % Adsorbed P
desorbed 0.3 1.4 1.1 1.3 5.8 2.5


Table 5. Effects of reaction time on phosphate adsorption
maximum and sorption energy constant for the
goethite system.


Sorption Energy
Reaction time Adsorption Maximum Constant
(hours) (Og P/g) (ml/Og P) 1/10 1340 0.66 1/2 2280 0.84 8 3850 2.40 18 to 76 4460 4.60


Table 6. The logarithm of (a) equilibrium P concentration
as a function of the logarithm of (b) amount of
P adsorbed.


log a

-.45 .15 3.07 5.54 log b 3.40 3.55 3.60 3.65









imaginary limit of the physically adsorbed ions and the outer solution containing all free ions. This could be expressed as:

(M) /n(H2P04)O (M)l/n(H P0) = k (89)
2 0 1 2 4
where (M) and (M)i are the concentrations of the cation M of valence n in the outer solution and inner solution respectively, terms (H2P04)0 and (H2P04)i are the phosphate concentrations in the outer and inner solutions, respectively, and k is a constant. From the above equation, it was evident that as salt concentration (M)0 increased, the (H2PO 4)� concentration must decrease. As a result, phosphate may be adsorbed through cation bridging or precipitation as cationphosphates.


Phosphate Desorption

Effects of phosphate free solutions of 1 M AlCl3, and

CaCl2 and NaCl are shown in Fig. 23. There was a significant effect of the type of salt on the extent of phosphate desorption. The magnitude of differences due to salt sources increased as the equilibrium phosphate concentration increased. For a low equilibrium concentration (1 pg P/ml), the desorption with 1 M AlCl3 solution released 550 og P/g of goethite. This was 3.1 and 1.8 times greater than similar desorption by 1 V CaCl2 and 1 1 NaCl, respectively. The higher desorption of phosphate by 1 M AlCl3 solution agreed with the apparent lower potential binding energy constant of





87





600



AlCl3



500 CaC12







400 0

0

to


300
0NaCI
.






200



150 100



so
0










0 1 2 3
EquilibriuIM Concentration (ug P/ml)
Fig. 23. Phosphate desorption in 10 hours from goethite
by three salt solutions at various equilibrium
P concentration.




Full Text

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INTERACTIONS OF ORTHOPHOSPHATE WITH IRON OXYHYDROXIDE MINERALS FOUND IN SOILS BY BOURAHIM YEKINI A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT OF THE REQRUIEMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY UNIVERSITY OF FLORIDA 1980

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The African People

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ACKNOWLEDGEMENTS The author wishes to express his sincere appreciation to Dr. John G. A. Fiskell, chairman of the supervisory committee, for his guidance and assistance throughout the entire course of this study, and for his valuable suggestions and excellent assistance in the preparation of this manuscript. The author is pleased to extend his sincere acknowledgements to Dr. J. J. Street and Dr. V. E. Berkheiser for their participation in the supervisory committee and constructive criticism of this manuscript. Appreciations are also extended to Dr. W. K. Robertson, Dr. R. C. Stoufer and Dr. T. L. Yuan for various assistances. Special recognition is expressed to Dr. Charles F. Eno, chairman. Soil Science Department, Dr. B. G. Volk, and Dr. D. F. Rothwell. Sincere appreciation is expressed to the AfricanAmerican Institute which sponsored this study. Appreciations are also extended to Carol Giles for her excellence in typing. The author wishes to express the deepest gratitude to his mother Mariama, his wife Olga and daughter Maryam, and his brothers and sisters for their moral support. Ill

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TABLE OF CONTENTS Page ACKNOWLEDGEMENTS iii ABSTRACT xi INTRODUCTION 1 CHAPTER I REVIEW OF SOME MODELS DESCRIBING FACTORS AFFECTING NUTRIENT AVAILABILITY 4 Nutrient Potential 4 Capacity and Intensity Relationship 7 Energy of Adsorption 8 Some Adsorption Isotherm Models 10 Effect of Surface Heterogeneity on Adsorption-Desorption 12 Generalization of Adsorption Isotherms 14 Soil Phosphorus Reaction Mechamisms 16 Specific Adsorption 19 Infrared Study of Phosphate Specific Adsorption 21 II KINETICS OF ADSORPTION AND DESORPTION 26 Adsorption 26 Desorption 27 Adsorption and Desorption Relationships .... 27 A Kinetic Model for Adsorption-Desorption ... 28 III THERMODYNAMICS OF ADSORPTIONDESORPTION REACTION 33 The Surface Charge 33 Zero Point of Charge (ZPC) 35 Thermodynamics of Adsorption 36 Free Energy as a Function of Distance ..... 36 Free Energy for Irreversible Fixation 38 Relationship of Surface Charge to Surface Potential 39 Surface Tension and Specific Adsorption .... 40 IV

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TABLE OF CONTENTS (Cont inued) CHAPTER Page IV MATERIALS AND METHODS 41 Geothite Preparation 41 Some Factors Affecting the Kinetics of Adsorption-Desorption 42 Surface Charge as Affected by Phosphate Adsorption 45 Soil pH, Iron Oxides, and Extractable P . . . . 46 V RESULTS AND DISCUSSION 48 Goethite and Phosphated Goethite Study by Infrared Spectroscopy 48 Goethite Structure Identification by Infrared 49 Factors Affecting Goethite Crystallization 60 Phosphate Adsorption and Desorption Studies 69 Effect of the Supporting Electrolyte 77 Time of Reaction Effects 88 Effects of pH on F Adsorption and Desorption 99 Effects of P Adsorption on Surface Charge of Goethite 108 VI SUMMARY AND CONCLUSION 112 LITERATURE CITED 118 BIOGRAPHICAL SKETCH 123 V

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LIST OF TABLES Table Page 1 Soil pH, iron oxide, and extractable P 47 2 Type of salt effect on the P adsorption on gothite at 5 mg P/ml 79 3 Effects of three electrolyte salts on the phosphate sorption maximum and sorption energy constant for goethite 81 4 Phosphate desorption from goethite by different anions 85 5 Effect of reaction time on phosphate adsorption maxim.um and sorption energy constant for the goethite system 85 6 The logarithm of (a) equilibrium P concentration as a function of the logarithm of (b) the amount P adsorbed 85 7 Effects of P adsorption on the equilibrium solution pH after 24 hours of adsorption reaction 101 VI

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LIST OF FIGURES Figure 1 Schematic representation of solidinterface-solution system 2 Infrared bands of (A) goethite; (B) phosphate added to goethite at the end of ageing; and (C) phosphate added to goethite at beginning of ageing. Curves A, B and C are for 22°C and curves A', B' and C are for 55°C 3 Infrared bands of (A) goethite; (B) phosphate added to goethite at end of ageing; and (C) phosphate added to goethite at beginning of ageing. Curves A, B and C are for 22®C and curves A', B' and C are for 55°C 4 Infrared bands of (A) goethite; (B) phosphate added to goethite at end of ageing; and (C) phosphate added to goethite at beginning of ageing. Curves A, B and C are for 22°C and curves A', B' and C are for 55°C 5 Infrared bands of goethite and lepidocrocite. (After Farmer and Palmieri 1975) . 6 Infrared bands of (A) goethite; (B) phosphate added to goethite at end ageing; and (C) phosphate added to goethite at beginning ageing. Curves A, B and C are. for 22°C and curves A', B' and C are for 55°C 7 Infrared bands of (A) Fe hydroxide material; (B) phosphate added to Fe hydroxide material at end ageing; and (C) phosphate added to Fe Hydroxide material at beginning ageing. Curves A, B and C are for 22°C and curves A', B' and C are for 55°C Page 29 50 51 54 55 56 57 Vll

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LIST OF FIGURES (Continued) Figure Page 8 X-ray diffraction patterns of (A) goethite, (B) phosphated goethite at end of ageing, and (C) phosphated goethite at beginning of ageing 59 9 Infrared bands of phosphated goethite at beginning of ageing for suspensions of various P/Fe values at OH/Fe = 6 62 10 Infrared bands of phosphated goethite at beginning of ageing for various P/Fe ratios at OH/Fe = 3.0 63 11 Infrared bands of phosphated goethite at beginning of ageing for various P/Fe ratios at OH/Fe = 1.5 64 12 Infrared bands of 1) goethite digested in D^O and 2) phosphated goethite digested in D^O 66 13 Infrared band of phosphated goethite after desorption by (1) 0.1 ^ KOI, (2) H„0, (3) 0.1 N KNOg, and (4) 0.1 NNa^SO^^ . 67 14 The 001 face of goethite lattice. (After Bragg and Claringbul 1965) 69 15 The 001 face of phosphated goethite 69 16 Adsorption isotherms as affected by the reaction times for goethite-solution systems 71 17 Adsorption isotherm for the Kenya soil after 24 hours of reaction time 72 18 Adsorption isotherm for the Georgia soil after a reaction time of 24 hours 73 19 Phosphate adsorption isotherm for the Colorado soil after 24 hours of reaction time 74 viii

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LIST OF FIGURES (Continued) Figure Page 20 Effects of the initial P concentration on the phosphate adsorption by goethitesolution (1 g/1000 ml) 76 21 Effects of the type of supporting electrolyte on P adsorption on goethite suspension (1 g/1000 ml) 78 22 Effects of the supporting electrolyte concentration on the kinetics of P adsorption by the goethite-solution system (1 g/1000 ml) 84 23 Phosphate desorption in 10 hours from goethite by three salt solutions at various equilibrium P concentration 87 24 Transformed Langmuir equations for phosphate sorption by goethite as affected by reaction times, a is equilibrium concentration and b amount adsorbed 80 25 Change in the equilibrium P concentration and P sorption as affected by the reaction time 89 26 Change in the equilibrium P concentration for three soils as affected by the reaction times 90 27 Effect of reaction time and equilibrium P concentration on P desorbed by 0.01 N CaCl_ from Kenya soil ~ 95 28 Effect of reaction time and equilibrium P concentration on P desorbed by 0.01 N CaCl^ from Georgia and Colorado soils 97 29 Effect on P desorbed by 0.5 N NHj^F as affected by the adsorption reaction time and equilibrium P concentration 98 IX

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LIST OF FIGURE (Continued) figure Page 30 Logarithmic plot of equilibrium P concentration change with time (t) relative to equilibrium at 18 hours (t ) for phosphated goethite ° 93 31 Effects of change in equilibrium solution pH on P adsorption and 0.01 CaCl desorption of P in goethite system. Initial P concentration is 5 pg P/ml 100 32 Effect of phosphate adsorption time on phosphate desorption by 0.01 ^ CaCl at pH 2 and pH 10 T 104 33 Effects of solution pH on the amount of P desorbed from goethite 106 34 Effects of pH and concentration of supporting electrolyte on phosphate desorption ...... 107 35 Potentiometric titration of goethite. Note zero point of charge occurs at pH 5.8 109 36 Potentiometric titration of phosphated goethite. Note zero point of charge occurs at pH 5.2 110 X

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Abstract of Dissertation Presented to the Graduate Council of the University of Florida in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy INTERACTIONS OF ORTHOPHOSPHATE WITH IRON OXYHYDROXIDE MINERALS FOUND IN SOILS By Bourahim Yekini June 1980 Chairman: Dr. John G. A. Fiskell Major Department: Soil Science The mechanisms, kinetics, and reversibility of orthophosphate adsorption on synthetic goethite and soils were investigated. Goethite was prepared by ageing of the precipiPsTs which appears after mixing PeCl 2 and NaOH solutions. Both the increase in Fe/OH ratio in the suspension and in ageing temperature favored a higher degree of goethite crystallization within a short period of time. With either OH/Fe = 6 in the suspension at room temperature or OH/Fe = 3 at 55°C and ageing for one week, good goethite yield was obtained. Infrared bands characteristic of goethite were at 3200 cm ^ (OH streching) and at 890 cm~^ and 790 cm~^ (both are Fe-OH bending vibration). The goethite structure was confirmed by X-ray diffraction intensities of the 4.19 8 and 2.70 8 peaks. The presence of phosphate at the beginning of goethite ageing weakened the goethite structure. I-Then F/Fe = 3.2 (regardless of OH/Fe ratio) the bond (Fe)0-P vibration at 1000 cm ^ was predominant, thus preventing the formation XI

PAGE 12

of Fe-OPI bond even after ageing of a suspension. When the addition of phosphate was made at the end of the goethite ageing, absorption of phosphate was determined as surface binding through binuclear bridging of the HPO^^ ion which was identified by the presence of bands at 1120, 1085, and 1030 cm The phosphate was assumed to penetrate the goethite structure whenever the vibrational band at 1000 cm~^ was present. Using synthetic goethite as supporting medium, it was found that the Langmuir and Freundlich equations could be used to describe the relation between the amount of P adsorbed and that remaining in the equilibrium solution for particular periods of reaction time. A kinetic model was proposed to describe the change with time of each of the following ion forms: 1) the free ions in solution, 2) the physically adsorbed, 3) the reversible chemically adsorbed ions, and 4) the irreversible chemically adsorbed ions . The amount of phosphate adsorbed increased with the increase in initial P concentration and the adsorption reaction time. The time required for the goethite phosphate systemi to reach an equilibrium state increased with the initial P concentration. Generally, the equilibrium state was reached after 18 hours for goethitesolution and 22 hours for soil-solution systems. The adsorption on goethite was increased by both multivalent cations XXI

PAGE 13

and the concentration of the supporting electrolyte. The potential binding energy constant increased (from 0.70 to 4.60 ml/ yg P) as the reaction time increased, but decreased as the cation valence of the electrolyte was increased. At the same initial P concentration, the amount of P adsorbed decreased almost linearly with change in pH according to the relation yg P/g goethite = -446 pH + 5680. The amount of P desorbed over a wide range pPI or supporting electrolyte remained nearly constant when the desorbing time was greater than or equal to 6 hours. The P desorbed from goethite and/or soils increased both with the initial equilibrium P concentration employed and the reaction time. At the same initial P adsorbed on goethite, the amount of P desorbed from goethite decreased when the pH was below 5.5 and increased when the pH was greater than 6 . Phosphate adsorption on goethite was found to induce a net increase of negative charge so that the zero point of charge declined from pH 5.8 to 5.2. Xlll

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INTRODUCTION Increase in urbanization makes it necessary to increase both the quantity and the quality of agricultural production in the tropical region. The implementation of an adequate agricultural policy involves improvements of the actual level of technology. This implies a clear understanding that the macroscopic relation be sought between the major factors producing food for human population through proper managements of soils, plants, animals, and insects under diverse climates and social activities. Modeling of the agricultural system is a useful tool for understanding a united approach for all major factors governing the system. With a suitable model, it should be possible to envisage needed change in a particular factor in order to give results as close as possible to a reasonable expectation. Only some microscopic relationships within the soils will retain our attention in this study. In a soil-solution system, the dominant soil phenomena taking place simultaneously are mass transport, diffusior^ adsorption-desorption , precipitationdissolution, and microbial immobilization and mineralization. The dominance of each soil phenomenon depends on the soil structure and texture, organic m.atter content, water content, temperature, and the ionic suite (types and concentrations). 1

PAGE 15

2 Upon fertilization, tropical soils may not react in the same way as do the temperate soils because of their different mineralogical , and chemical properties. Among the major elements, phosphorus is, after nitrogen, the most deficient nutrient and its availability is strongly dependent on the mineralogical composition of the soil (type of clay, and metal oxides). The low availability of phosphate to tropical plants is due both to the high soil phosphate fixing capacity and to precipitation of fertilizer phosphorus from the soil solution. Not all forms of phosphorus bound on the oxide surfaces are held with the same degree of strength. V/ith time, as phosphorus is depleted from the solution, some phosphate may be replenished by release from the solid phase. The magnitude of this release, vjith respect to time, is dependent on soil characteristics. The initial concentration in solution is not a sufficient measure of phosphorus availability. In tropical soils, iron oxides are quite important in determinig the solid phase capacity to fix and supply phosphorus to the soil solution. How some factors, such as OH/Fe ratio, temperature, and phosphate concentration affect the formation of goethite are herein investigated. Infrared spectroscopy will he used to identify the effects of soluble phosphate on goethite structure and the nature of phosphate binding on goethite. Synthetic goethite as well as three soils will be used to investigate some

PAGE 16

3 factors affecting the time dependence and the degree of reversibility of orthophosphate adsorption.

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CHAPTER I REVIEV/ OE SOME MODELS DESCRIBING EACTORS AFFECTING NUTRIENT AVAILABILITY Nutrient Potential The concept that availability of a nutrient can be described by its potential activity was first brought to general notice by Schofield (1955). His concept was that in comparing two soils of different water holding capacity which contain the same amount of available water, the one having the lower capacity has a lower potential and the water is easier to extract than from the soil of higher holding capacity. Their pF value defines the energy v;ith which water is held on the soil particle surface. The potential given by the pE value provides the basis for quantitative evaluation of the availability of water. As water is taken out of a soil, it is replaced by lateral movement of groundwater. Similarly, as a nutrient is taken from the soil solution, it is replaced by other ions, by desorption, or by diffusion as well as mass flow. Eventually, it is not the amount of a nutrient in a soil that primarily controls the uptake of that nutrient by the plant, but the work required to withdraw it from the solution. This work may be related to Gibb's free energy (G^) if the uptake is mainly from the soil solution. To derive an expression for 4

PAGE 18

5 the nutrient potential, three postulates have been advanced by different workers. Postulate 1 Where the solution is at constant temperature and pressure, and when the system is at equilibrium, is uniform throughout the soil solution so that G. = G? + RTLn a^ + zF'l'f ( 1 ) where G? is the standard molar free energy of an ion i in the solution relative to an arbitrarily zero of the electric potential, a^ is the activity of this ion in solution, z is the valence, F is the Faraday constant, and is the electric field effects. Postulate 2 When a soil is in equilibrium with a solution, the electro— chemical potential of the ion is constant throughout the system (soil-solution). Then, the partial molar free energy of any ion in the soil complex can be determined by analyzing the solution. This free energy of the ion in solution is assumed to be constant even if the solution is separated from the solid phase by centrifugation. As a result, if the ion is taken out of the field influence, then zF4' = 0 where z, F, and 4' are as described in the previous section. Then the relationship in Eq . (1) can be reduced to

PAGE 19

6 G. = G? + RTLn af (2) Postulate 3 This postulate is that potentials are determined by transfers of ions. In order to maintain the electroneutrality, equal transfer must occur for quantities of positive and negative ions in one direction or, alternatively, transfer of one ion in the first direction occurs as an equivalent quantity of ions of the same sign moves in the opposite direction (Barrow et al. 1965). It is also assumed that divalent cations (Ca and Mg) dominate m.ost soil surfaces as exchangeable cations and in the solution, except for saline soils. For a situation in which transfer of Ca^"*" is in one direction and KÂ’*Â’ in the other, the net change in free energy is ^^Ca,K " ^^Ca,K ^Ca ^ For phosphorus, over the range of pH v;hich exists in soil, two forms of the orthophosphate ions are in equilibrium; thus H PO" + H O ^ = = = = = = i HPO^^~ + H.O'^ (4) pK=7.2 ^ Assuming that plants absorb m.ainly H^PO^ ion and not HPO^~, then a cation such as Ca must accompany the H^PO^. The equation for the transfer of both ions from the equilibrium solution is

PAGE 20

7 ^^Ca(H 2 PO ^)2 ^ RT(l/2pCa + pH^PO^^) (5) A common use of the relationship is the Schofield's phosphorus potential, where I is given by I = 1/2 Ca^"*" + pH^PO^. Capacity and Intensity Relationship Phosphorus ions pass from one phase to another as a result of chemical potential differences between the phases. In general, any substance tends to pass from a region of higher chemical potential to a lower one. The ability of the solid phase to supply phosphorus to the solution phase, as it is depleted, can be termed the capacity. For any ion such as phosphorus, at equilibrium, the phosphorus buffering capacity (PBC) is a soil characteristic where PBC is dAP/dl. The term with I = l/2pCa^'^ + pH^PO" and AP is the gain or loss in phosphorus by the soil. The relationship between capacity and intensity factors is expressed as Q/I curves which are generally composed of two parts with linear and curvilinear portions. In the Q/I relationship for potassium, the curved portion is attributed to a number of sites which have specific affinity for potassium at lov; concentration and the linear part is associated to non-specific sites, Beckett (1971). Ke also believed that, in general, the curved part of the Q/I relation can be represented by a Langmuir adsorption isotherm whereas the linear part fits the Gapon relationship.

PAGE 21

8 Becketrlr (1971) affirTned "that "the non-linear part of the curve represents a certain region where there is a definite limited number of sites on the exchangeable surfaces and exhibits a selective binding power for the adsorbing ion: The Langmuir adsorption isotherm could be used to describe such a situation. In addition, if Eq. (3) is recalled, the activity ratio (AR) is K'^/CCa)^^^ of the ions in solution which is the product of the concentration of the exchangeable ions and the Gapon constant (k^). Energy of Adsorption Partition function To be effective, the collision between molecules and the collision surface must provide a certain m.inimum amount of energy called the activation energy. Since the activated complex is a transitory species, the equilibrium constant cannot be measured experimentally. However, the partition function arises from the quantum theory that a molecule can exist only in states with definite energy limits. from Boltzman distribution, it is recognized that where is the total number of molecules, N . is the number o/\ of molecules at zero energy, E^ is the activation energy at the state i , and g^ is a constant. In this case, also quantity Q = Zg^^^exp ( -E^^/kT ) which is the partition function.

PAGE 22

9 Electric Field Effect The following treatment was first derived for solid-gas interaction and is known as Thin's theory reported by Brunauer (1943). The strength of the electric field surrounding the adsorbent is given by E — 9 where e^^ is the charge distribution of the absorbent, e is the dielectric constant of the gas, and r is the distance from the surface . The force (F) acting on an induced dipole is F = ke^v/4 nr^^"^^ where v is a constant and k = e-1/ e . The adsorption potenTis-l is the force multiplied by the distance through which its acts (F.Ar). Heat of Adsorption Assuming no work due to phase change is done during the adsorption-desorption, the molar energy of the ion in free solution (u^) and the energy of the ion in the adsorbed phase (Ug) are related to the loss of ion from the solution due to adsorption as shov/n by AH = b(u^ vjhere b is the number of miole adsorbed, and AH is the integral heat of adsorption. Differential Heat of Adsorption Assuming that the adsorption process is reversible AH can be obtained from data at two temperatures (T^ and T^ ) by

PAGE 23

10 using the Clausius-Clapeyron equation for a constant surface coverage, namely ^2 AH 1 1 R 1 1 ro 1 where a^ and a^ are equilibrium concentrations at temperature and respectively, and R is the molar gas constant Some Adsorption Isotherm Models Freundlich Equation The Freundlich equation was first introduced in an empirical form (Bach and V/illiams 1971). It assumed that the energy of adsorption decreases exponentially with increas ing saturation of the surface. The equation is b/m = ka^^^ ( 9 ) where b/m is the amount of phosphorus adsorbed per unit weight of soil at the equilibrium concentration termed a, and where k and n are constants. The logarithmic form of Eq. (9) is log(b/m) = log k + 1/n log a (10) The plot of log(b/m) versus log a should yield a straight line. This equation is valid only in a limited range of concentrations. By taking into account the initially exchangeable phosphorus and the ability of the soil to lose or gain phosphorus (AP) during the equilibrium reaction, the following modification was introduced AP = Aa = ka^'^Â’^e (11)

PAGE 24

11 where Aa is the amount of phosphorus gained or lost, e is the phosphorus initially present, and n is a constant. Modified Freundlich Equation In a colloid-solution system, the adsorption is timedependent. Kuo and Lotse (1972) introduced a time factor * }c t 1-e 2 after assuming that the constant is small. They found that b/m = ka^t^'^^ (12) where a^ is the initial phosphorus concentration. Their study indicated that the rate constant increased with an increase in concentration. Langmuir Equation The derivation of the Langmuir equation is based on the following assumptions: 1) the energy of adsorption is constant and independent of the degree of coverage, 2) there is not interaction between adjacent adsorbed molecules on an homogenous surface. If the system is in dynamic equilibrium, this results when the rate of adsorption is equal to the rate of molecules escaping from sorption surface, so that ^ 1 ® ^2 ^ 13 ) where and k^ are the rate constants of adsorption and desorption, P is the gas pressure, and 6 is the fraction of the surface coverage by the gas. By rearrangement, the relationship is

PAGE 25

12 (14a) and by analogy for soil-solution system, where a is the ion concentration in solution. The above equation can be put in the linear form as follows where k is the adsorption energy constant and s is the adsorption maximum. Different workers have observed that there is not a linearity between a/b/m and a. This may be caused by at least two types of sites of adsorption having different energies of adsorption. To fit this criterion, the Langmuir equation was modified to be where k^, s^ are constants for region 1, and k^, S 2 are constants for region 2 (Syers et al. 1973). Effect of Surface Heterogeneity on Adsorption-Desorption Langmuir and Freundlich equations can be obtained from Toth's equation (Jossen et al. 1978), for homogeneous surface. Toth's equation is obtained by integrating the equation a/b/m = a/s + 1/ks (14c) (14d) ( 15 ) Where 6=0, the Freundlich equation is obtained. Where 6=1, the Langmuir type of equation is

PAGE 26

13 ( 16 ) where 6 = d exp(-E/RT), a is the equilibrium concentration, Another factor is that the energy of activation (E) is constant for a homogeneous surface but varies on a heterogeneous surface with the degree of coverage (Jossen et al. 1978). Assuming that the free energy of activation for adsorption varies linearly with coverage of the surface, Lingstom et al. (1970) proposed the following model where A is the adsorbate, S is the surface, and and k 2 are the rate constants. They deduced that the rate of adsorption process is de/dt = k^(l-0)(l-l/26)e + k 2 [(l-l/ 20 )e ®^-l/20^e^®] (18) Another useful relationship is the Elovich equation where the activation energy is a linear function of the amount adsorbed Aharoni and Ungarish (1977) modified this equation by introducing the fact that the heterogeneous surface is comprised of a large number of homogeneous regions having unequal number of adsorption sites. b the amount adsorbed and b is the maximum adsorption. (17) where 0 = b/b°°. E = E + ab o (19a)

PAGE 27

14 = Eq + RTln(gy^/n^ + y) (19b) where = db/dE, g and y ai'e constant, and n^ is the number of sites of energy , where E^ is the activation energy characteristic of a region. They also assumed that, at any moment, adsorption takes place preferentially on the region that has the lowest activation energy at that moment. Considering the case where equilibrium is attained at a region when y^ = y^^j then it follows that y^^ = k^exp(E/RT). The overall rate of adsorption is given by db/dt = k N exp(-E./RT) and N. = /n„dE (20) o t ^ t t E where is the number of sites with an energy E^ at time t and n^ is the number of sites with an activation energy between E and E + dE. Generalization of Adsorption Isotherms Freundlich and Langmuir equations can be extended for cases of competitive adsorption (Jaroniec and Toth 1976; Digiand et al. 1978). In this case, the Freundlich type is given by b^ = k(Za^) (21) i and the Langmuir type by b^ = ka^/ (k^ + k^a^ + ^2^2 (22) which is a partial isotherm of a^ relative to the total ions present. A binary system can be simplified if the following

PAGE 28

15 assumptions are made, 1) the surface is formed by a collecregions each being characterized by 3. binding energy (E), 2) the number of surface atoms Ih with energy E follow the Boltzman distribution, which is described by Ni = N^exp(-e/RT) (23a) and 3) all adsorbed molecules stay on the surface for the same average length of time 6 so that 6 = 6^exp(E/RT) (23b) During the time 6, the fluctuation of the number of adsorbed and desorbed molecules compensate between themselves randomly (Vlad and Segal 1979). They considered that the adsorption energy is an increasing function of the extra energy e. From Mclaurin development, the extra energy can be expressed as e = Za^(E E )^ (24) where E^ is the smallest adsorption energy corresponding to the value of zero energy and n is generally integers of 1 or 2 . The energy distribution X(e) is expressed by Van Dongen approximation as X(e) = exp(Z a^e) i = 0, 1, or 2. (25a) i or generally as N , N X(e) = 1/kT I na (E E )^~-^exp[-l/kTZa (E E )^] n = T ^ ^ m (25b)

PAGE 29

16 The general isotherm is expressed as b(a,T) = /^b^(a,e) X(s) de (26) where b(a,T) is the fraction of the total surface covered at (a,e), b^(a,e) is the local isotherm which may be analogous to an isotherm equation for an homogeneous surface or region, is the range of possible variation of adsorption energy assuming that = [0,«>]. Temperature Dependence of the Rate Constant According to the Arrhenius equation, it is known that ^ads " Aexp(-E^/RT) (27a) where is the energy of activation, R is the gas constant, and T is the absolute temperature. Erom the transition state theory, this relationship is extended to kads = (kT/h)exp(AS^/R)exp(-AH^/RT) (27b) = (kTZh)exp(AG /RT) (27c) where AS is the entropy of activation, AH = E RT which is the enthalpy of activation, and is the Gibb's free energy of activation. Soil Phosphorus Reaction Mechanisms Plant responses to soil phosphorus are a function of the solubility of phosphorus. Any factor altering the solubility of phosphorus will also alter the plant response

PAGE 30

17 (Hemwall 1957). The solubility of phosphorus depends on several factors such as the composition of the soil solution, the pH of the solution, and temperature of the soilsolution system. Calcareous and Neutral Soils The adsorption of phosphorus on calcareous surfaces can take place by replacement of water molecules, bicarbonate, and certain other anions or cations. The relative strength of the phosphate ion adsorption depends on the solubility of the compound at the calcium surface (Kuo and Lotse 1972). The precipitation of phosphorus may be due to the formation of a whole series of insoluble calcium phosphates. Some of these with associated solubility products expressed as the pK value are reported by Lindsay and Moreno (1960). Compounds Chemical formula pK Calcium phosphate anhydride CaHPO^^ 6.66 Calcium phosphate dihydrate CaHPO^^. 2 H 2 O 6 . 56 Octocalcium phosphate Ca^^H(P0^^)3. SH^O 46.91 Hydroxyapatite Caio(PO^^)6.(OH)2 113 .70 Fluorapatite Caio(PO^^)6.p2 118.40 In calcareous soil, hydroxyapatite and fluorapatite are the major phosphorus compounds, whereas in neutral soil, octocalcium phosphate becomes important and in moderately acidic soils dicalcium phosphate may occur.

PAGE 31

18 Acid Soils Iron, aluminum and pH are the main factors controlling phosphorus solubilities; In acid soils, the solubility of phosphorus increases with decreasing free iron and aluminum activities and with increasing pH. The form of phosphate sorbed by acidic soils was recovered (90%) as iron and aluminum phosphates (Ghani and Islam, 1946). Yuan et al. (1960) reported that up to 80% of the added phosphorus to acid sandy soils was present as aluminum phosphate and 10% as iron phosphate, but when the reaction temperature was increased more iron phosphate was formed. After reaction of phosphorus with soluble Al^'*', a microscopic examination showed an hexagonally shaped crystal in which the interplanar spacing was similar to those of palmerite (Haseman et al. 1950). In general, some of the main forms in which phosphorus can precipitate with iron and aluminum are given below. Compounds Chemical formula pK 30.5 Variscite AlPO^^. 2 H 2 O Strengite FePO^^. 2 H 2 O 34.3 Taranakite (K,NH^^)3HgAl^(PO^^)g.l8H2 0 176 . 0 Palmerite HK2Al2(PO(^)g. 7H^0 Since Al and Fe atoms are part of the surface colloids, they react with phosphates. Whether the process is

PAGE 32

19 precipitation or adsorption depends on the size of the metal polymer and the pH of the phosphate, and its concentra tion. In a moderately acidic medium, with a high phosphorus concentration, the reaction process may be typically precipi tation, resembling that reported by Hsu (1965), [Alg(0H)^2^^^ + 6 H^PO” Alg(OH)^2^H2PO^^)g (28) Specific Adsorption General Description In a solid-solution system, adsorption of molecules (or ions) occurs when there is a change in phase from the free state (in solution) to the bound state at the interface. Ions deposited at the surface likely orient to form the Stern layer as some ions may approach closely to the surface structure. In this case, the ions are said to be specifically adsorbed on the surface. Non-specifically adsorbed ions are either in the diffuse Gouy-Chapman region separated from the solid surface by at least one molecule, or electrostatically bound to that surface. Specifically adsorbed ions are in the coordination shell of the surface atoms and are maintained there through chemical binding (covalent or coordinate binding). Since specific adsorption of cations or anions occurs even when the surface possesses a net positive or negative charge, respectively, there must be an electrostatic contribution due to polarization of the ion or molecule.

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20 Both chemical and electrostatic attractions contribute to the energy of adsorption. The magnitute of this energy of adsorption determines the degree of reversibility and the amount adsorbed. Since the reaction is pPI dependent, at any pH, there is a maximum adsorption of anions and when these maxima are plotted against pH, the curve is termed the adsorption envelope which is described by Kingston et al. (1967) as where A is the amount of the ion adsorbed per unit weight of adsorption, a is the degree of dissocation of acid anion, K is the dissociation constant of the most highly charged anion that is adsorbed, and V is the amount of ion adsorbed at the maximum level. Why som.e of the adsorbed anions such as phosphate may be irreversibly held is explained by certain postulated mechanisms of phosphate adsorption. The removal of phosphate from free solution is assumed to proceed through the replacement of coordinated H^O groups and/or some OH” ions. A = 2Vn/U a) = 2V Kun (29) (K + [H])^ OH OH 0-P = 0 M OH + OH (30) OH OH OH where M is a metal (Fe or Al ) , and bonded by a coordination link .

PAGE 34

21 Infra-red Study of Phosphate Specific Adsorption Theory of Infra-Red Spectroscopy If a molecule is placed in an electromagnetic field, a transfer of energy from the field to the molecule occurs when Bohr's frequency condition is satisfied: AE = hv, where AE is the difference in energy between two quantized states, h is Plank's constant, and v is the frequency of the light wave. Pure rotational vibrations are usually observed in the microwave vibrational spectra in the infrared whereas electronic spectra are in the visible and ultraviolet. The infrared spectra originate in the transition state existing between two vibrational levels of the molecules. From a classical point of view, a vibration is active in the infrared spectrum if the dipole moment of the molecule is changed during the vibration. The dipole moment Q is related to the strength of the electric field by Q = aE and Q is a vector whose direction is the line between the center of gravity of the protons and electrons . Let a diatomic molecule be represented by two masses m^ and m^ moving along the molecule's axis with displacements of and respectively. T'he displacements of the two atoms are induced by forces which can be obtained through Hooke's and Newton's law. KCX^ X^) = m^ dt dt 2 2 ( 31 )

PAGE 35

22 The solutions of these equations of motion are: = Acos(27TVt + a) and Acos(2iTVt + a) (31b) After differentiating and substituting back in Eq . (31), it can be found when solving for v (Nakamoto 1978; Colthup et al. 1975) (32) where k is the force constant and v is the frequency of vibration. The experimental observations of frequency of vibration of a crystal are due to bond vibrations of atoms within the unit cell. Such vibration m.ay help us to determine the nature of the atoms composing the unit cell by comparison to known polyatomic vibrations. Because of the interactions, the symmetry of a molecule is generally lower in the crystal. q line state than in the isolated state. The isolated PO^ ion is of tetrahedral form (Nakamoto 1978): 0 / t 0 The above tetrahedral structures predict two infrared active fundamentals, one is stretching (v^) and the other (^^) is bending .

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23 Infrared Identification of the Forms of P Adsorption From their studies, Kingston et al. (1974) proposed the reaction '^0H2 -2 0 (33) Reversibly adsorbed P Irreversibly adsorbed P Parfitt et al. (1975) identified a form of phosphorus which can be considered to be completely within the goethite structure, since the phosphate is bound with two Fe atoms and is H-bonded to a third Fe. Fe 0 HO 0 0 Fe 0 Fe Fe OH 0 or > 0 .0 Fe '0 Fe (34) Such a chemical binding of phosphate on the goethite surface is said to be specifically adsorbed, and can be considered as the formation of a new solid phase or growth of the solid phase . Colloid or Soil Surface Effects Substances dissolved in soil solution can move by molecular or ionic diffusion resulting from a concentration

PAGE 37

24 gradient within the solution, or by mass flow of the soil solution. This is complemented by sorption of the ions onto the soil surface, or by a combination of these factors. The above phenomena occur simultaneously and are described by the following equation of de Camargo et al. (1979): dt ^^2 dx dt dx (35) where a is the ion concentration in the soil solution, b is the amount of the ion adsorbed, x is the distance of the ion from the electric potential effect due to the surface charge on the charged solute. In Eq . (35), the constants are A which is a coefficient combining diffusion and hydrodynamic dispersion (cm /hour), B is the average pore density derived from the volumetric water content and C is the average pore water velocity, and D is the ratio between soil bulk density and volumetric water content. In a soil system where the terms Bda/dx and Ddf/dx approach zero in a soil column of a semi-infinite length, over a small time t, Eq. (35) can have the following solution given by Lingstrom et al. (1968): a(x,t) = a erfc X 2Kt (36) where a^ is the initial concentration, and K is a constant depending on the free energy of adsorption.

PAGE 38

25 Root Effects The effect of roots in a soil solution arises because ions movements occur into and out of the roots, with the net balance being an influx. This uptake is mainly associated with ion transport by diffusion and mass flow and is proportional to the concentration at the root surface. It is assumed that the rate of plant uptake is equal to the rate of loss of the solute, so that the reaction is: da dt = -6a (37) where t is the reaction time, and 6 is a constant. A solution of the above differential equation is given by Baldwin et al. (1973) in the following form: a aoexp -wt/(w, + W2X In ^) (38) X -r where a^ is the initial concentration in solution, and where w, w^, and w^ are constants, r is the root radiu,, and x is the distance from the root surface. This equation describes the change of the solution concentration at a certain distance from the root surface with time.

PAGE 39

CHAPTER II KINETICS OF ADSORPTION AND DESORPTION The interfacial region between soil colloid and solution is a center of intensive chemical and physical activities. The type of activities retaining our attention here is the adsorption and desorption processes of ions. Adsorption The adsorption of an ion on the colloid surface can be considered as a second order reaction involving an ion (A) and the sites (S) on the surface: ^1 A + S i AS (39) where is the rate constant of adsorption. By assuming that the reacting sites on the colloid surface are reacting species equivalent to (A), then ^ = k^[A]^ (40a) and rearranging this is = /^k, dt (40b) [A]^ t=0 ^ or can be written as — — = k,t (41) [A] [A]^ The plot of 1/[A] versus time t should yield a straight line. 26

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27 Desorption The desorption of any adsorbed ion (AS) can be considered as a pseudo-first order reaction: A + S where ]<2 is the tiation of Eq. rate constant of desorption. By differen(42), the following can be obtained: k^CA] (43) And by integrating the above differential equation, this becomes [A] = [A] exp(k„t) (44) o 2 where tA]^ is the initial concentration of the ion specie A, t is the reaction time, and k 2 is the rate constant of desorption . Adsorption and Desorption Relationships Since the adsorption and desorption phenomena are simultaneous a combination of the Eq . (41) and Eq . (44) is formed as follows : A + S ==P-== AS (45) ^2 d[A]/dt = -k^[A]^ (46) d[A]/dt = k2[AS] where A, S, t, k^ , and k 2 are described in the precedent ( 47 )

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28 sections. By dividing the Eq . (46) and Eq . (47), and rearranging, it can be obtained: d[A] _ ^ d[A] dCA]^ ^2 [A] 1/[A] 1/[A] = (k,/k„)ln([AS]/[AS] (48) o 1 2 o If we set a = 1/[A] 1/[A] and b = [AS], then Eq . (48) can o ^ be written as ^2 b = b exp(:r— a ) ( 49 ) o where b is the amount initially adsorbed. The above equations conbining adsorption and desorption phenomena are valid only at equilibrium, there is need to propose another model combining adsorption and desorption but valid at any reaction time. A Kinetic Model for Adsorption-Desorption Energy Constant Characteristic of the Phase For a collision to be effective in producing molecular species from reactants, some amount of energy must be available to allow for the necessary bonds to break and be formed. In 1889, Arrhenius suggested that molecules must get into an activated state before they become reactive. In any system, an equilibrium exists between ordinary and active molecules and only the latter are rich enough energetically to undergo reaction

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29 A ==== (A-B)^ === B (50) # where (A-B) is the spatial configuration of the transition state. The species (A-B) does not represent the active molecule of Arrhenius. His active or energized species are rather a few reactant molecules having sufficient energy to • # get into the transition state, (A-B) , or activated complex condition, but not necessarily having the spatial configuration corresponding to the transition state (Eying and Eying 1963). Let us consider the following system /Timiin / — //// C III B ///I//////// 7 ! / //I/////// 1 //I/////// 1 1 1/1111111' / ///////// / / Figure 1. Schematic representation of solid-interfacesolution system. where A is the solution phase, B is the interfacial phase, and C is the solid phase. Here A, B, and C are different degrees of energized species of the same ion (or molecule) with respect to the solid surface; these are considered characteristic of solution, interfacial, and solid phase, respectively. The surface sites at which ions take the

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30 form C are at random on the solid surface. It is believed that the sites associated with the lowest energy are the first to be filled by the ions taking the form C. Formulation of the Model In the system described above, it can be assumed that the following type of reaction is taking place: A ^-1 k2 B, = C b k_2 c (51) where A, B, and C are as described previously and subscripts a, b, and c designate the amounts of the ion in the forms A, B, and C per unit volume of the system. The terms k^ k 2 are the rate constants for forward reaction and reverse reactions, correspondingly are k-1 and k-2. The rate of a reaction can be expressed by the rate of change of the concentration of the ion species: da , , , dt ^-1^ ^1^ (52a) which in expanded form becomes dt ^ ^1^ k_ib + k_2C -k2b = k^a (k_^+k2)b + k_2C (52b) and dc dt k2b -k_2C (52c) Since any colloid cannot, within a finite period of time, indefinitely fix an ion in the C energy form, then c must reach a maximum s at equilibrium. In Eq . (52b), it is assumed that k _2 is very low compared to k 2 .

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31 The simultaneous solution of the differential for Eq. (52a) and Eq. (52b) is w,t w„t a = ae and b = Be (53) where w^t W^t Wj^t aw^e = -k^ae + k_^Be (54a) and w,t W-|t w„t av]^e = k^ae (k_^ + K 2 ) Be (54b) From the Eq. (54a) and Eq. (54b), the coefficients a and B can be determined as a function of the rate constants. In any soilsolution system, equilibrium is attained when the rate of change of each ion in the different phase becomes nil, so that it is then da dt 0 such that k 2 b k_ 2 C duced as dc dt 0 0 and c=s, then the term B is deSimilarly , value, a i „ ’'-2 6 = if k ^b k^a s found to be (55) 0, then substituting in the B a s e -w. t o In the initial a and b expressions, shown in Eq . (SB) (53) , it can be shown that

PAGE 45

32 and k , k „ w, (t-t ) ^ ^ -1 -2 1 o ^ ^ k — ic — ® k „ w^(t-t^) W — ^ ^ ^ b = s — e ^2 (57a) (57b) The Eq. (52c) can be expressed in the following as a first order linear differential equation: j w„t d| " ^-2^ = e^2^ (58) The solution to this differential equation to satisfy the condition where c = 0 at t = o can be written in the form 6k, c = k^+w^ ”2'' k ,t e e -2 (59a) If we assume that k _2 is negligible, c can be further simplif ied 6k c = k^+w^ 2 " 2 '' e (59b) where w^ and w^ are constants.

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CHAPTER III THERMODYNAMICS OF ADSORPTION-DESORPTION REACTION At the microregion in the colloid or soil solution extending from the surface to the outer limit of the first adsorbed layer, it is assumed that the electrical effect due to surface charge is negligible to avoid the difficulty in estimation of the electric field effect. In this microregion, there is ion interaction between both the surface and the electric potential effects (which are induced by the surface charge ) . The Surface Charge The colloidal behavior depends on how the surface charge originated. There are two type of colloidal charges. The first type has the charges due to crystalline imperfec4+ 3 + txon, such as isomorphic substitution of Si by A1 or by 2 + other cations such as (Mg ) having a lower charge. This type of charge is found in the crystal lattice of clays such as montmorillonite and vermiculite. In these cases, the charge density is constant per unit of surface area. The second type occurs where the surface charge may be created by preferential adsorption of a certain ion, such as that by hydroxyl or phosphate ions. In this case, the charge arises at the exterior edge of crystals or lattice, thereby inducing a constant surface potential (van Olphen 1977; Bolt 1976). 33

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34 Gouy-Chapman and Stern Theories of the Double-Layer From Gouy-Chapman theory, the charge density can be deduced considering the electroneutrality condition (van Olphen 1977; Sennet and Olivier 1965): where p is the space charge density (or net sum of positive and negative ion concentration, o is the surface charge, and is the surface potential. The Gouy-Chapman theory gives an over-estimation of the double layer capacity, namely where K is the reciprocal of the double layer thickness, a is the ion concentration, e is the electron charge, z is the valence of the ion, and e is the dielectric constant. Stern recognized the importance of the ion sizes near the surface. He proposed that the counter ions could be divided between a diffuse layer and an immobile surface layer of thickness 6 able to contain a maximum number of counter ions per unit of surface area. This may be expressed as CO 0 = -/^pdx (60) a K = (8z e a/e kT) (62) ( 63 ) where o is the charge corresponding to a monolayer of m

PAGE 48

35 coun'tGr’ions ) is "the van der Wall Gnengy, A is ihG frGquGncy factor, 1'6 is the GlGctric potantial at thG StGrn Layar. Zgpo Point of ChargG (ZPC) On a niGtal oxida surfaca, chargas ara creatad by tha adsorption and dasorption of or 0H“ ions which are affected by their concentrations in solution. The Parks and da Bruyn (1962) model is (64) where M is A1 or Fa with A and C are the associate anion and cation, respectively. From the Eq . (64), it can be seen That there is a pH at which the net surface charge is zeroj this point is termed the zero point of charge (ZPC). For metal oxides, such as descriWd above, H"*” and 0H~ ions are potential determining ions . By approximation to the Hernst equation, Keng and Uehara (1974) reported that .+ RT T H o (65) where R is the gas constant, T is the absolute temperature, F is the Faraday constant, and <<25 mV; the Gouy-Chapman double layer equation is reduced according to

PAGE 49

36 % = = W (0-059)(pH^ pH) (66) Thermodynamics of Adsorption Effect of the Electric Field In this system introduced in Fig. (1), the compartments A, B, and C can be considered as being three phases of a particular ion at different potential energy levels. In each phase, molecules are at different energy levels and varying within a range of potential characteristics of each phase . Around a colloid center, the ions exhibit a Boltzman type distribution: c = b exp [-F(T^ 'F^)/RT] (67a) b = a exp C-F(Tj^ T^)/RT] (67b) Supposing that 1' = 0, it can be shown that a ''o = '•'o = r I where T is the surface potential, and T are electric o ^ b a potentials that is characteristic of the phases B and A, respectively . Free Energy as a Function of Distance From the outer limit of the Stern layer, both Gaussian and Gouy-Chapman concepts can be applied to a particular ion. The movement of an ion (or molecule) results from a succession of collisions that may move it at random in a

PAGE 50

37 positive or negative direction. These ionic concentrations follow the Gaussian distribution: a = 7 » exp(-x^/4Dt) (68) 2(nDt) ' where M is the amount of substance deposited at the plane when X = 0 at the time t = 0, and D is a constant. Assuming that the Gouy-Chapman distribution holds in the following equation : T = exp(-Kx) (69) where the free energy can be expressed as G. ^ = G° + zFT + RTln a (70a) and * zFT^exp(-Kx) + RTln — At constant time t, this expression becomes RT dG ,, -Kx T— = -zFT Ke dx o 2W X (71) The limit between the physically adsorbed layer and the completely free ion in the phase A may be defined when dG/dx=0. If we set y = -Kx and y = -RT/ 2Dt/zFT^K, then the expression can be written y = 0 ( 72 ) and become y^ + 2y(l h + 1 = 0 where

PAGE 51

38 (e^ = 1 + y + 1/2 + . . . ) V7hen this quadratic equation gives two roots , they may be indicators of the transition states between different degrees of adsorption. Free Energy for Irreversible Fixation The active transport of an ionic substance against the gradient potential is determined by the difference between the total potential within each phase (A, B, and C). The maximum energy, other than expansion work for a change of ion activities, is expressed by at constant temperature and pressure. The transfer of ions between the phase B and C is due to a potential energy minimum for phase C written c B as G. and maximum for phase B written as G. . In this 1 ,m ^ 1 ,m case g 9 = z.F4'^ + X. (73) 1 ,m 1 1 and, correspondingly, G? = z.FH'® + RTln b (74) 1 ,m 1 where G^ ^ is also the Gibbs free energy of the ion i per unit mole, is defined as the minimum energy for irreversible fixation (including only the electro-chemical energy), B C F is the Faraday constant, T and T are the electric potential in the phases B and C, and the term RTln b is the chemical potential. The transfer of ions between the phases B and C is governed by the difference in total energy between the two phases.

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39 k_2 AG. = -46az.F + RTln s ^ + RTw„(t t ) X. (75) 1 ^2 z o 1 where 'i' = = 4II6a which is the Gaussian equation for a molecular condenser. This by substitution gives 2 AG. ^ = 4n6az.F + RTln s t— ^ X. (76) i,m 1 )<2 1 When the conditions are at equilibrium, or AG = 0 and t = t^, then k X. = 4n6az.F + FTln s (77) 1 1 K 2 It is assumed that the minimum energy X^ is independent of the state of equilibrium, then for a given colloid surface which has a particular surface charge, the free energy change AG. is reduced to 1 , m AG. ^ = RTw„(t t^) (78) 1 , m / o Relationship of Surface Charge to Surface Potential During the overall transfer of ions from the phase A to phase C at equilibrium conditions, after substituting the X^ value of Eq. (77) and assuming H' =0, the relationship becomes d>F o 446 a . F da 1 where 6 is the Stern layer thickness, F is the Faraday constant, and a^. is o the (79) is the surface charge, ion valence.

PAGE 53

40 Surface Tension and Specific Adsorption In the interfacial colloid-liquid region, some ions may enter into the coordination shell of surface atoms. As a result, there is a modification of the colloid surface (A) (expansion or, contraction) . This work required to expand or contract the surface when divided by the change in surface area is the surface tension (y) (Atkins 1978). following series demonstrates the relationship: The g 9 = z.FTc + Ay (80a) g 9 = z.FTc + X.c 11 1 Differentiating the Eq. (80a) and Eq . (80b) gives (80b) dG^*“ = z^FTdc + z^FcdT + ydA + Ady (81a) dG9 = z.FTdc + z.FcdT + X . dc + cdX . 11 1 1 1 However, it is known that (81b) dG9 = z.F'Fdc + X.dc 11 1 (82a) dG9 = z.FH'dc + ydA 1 1 ' (82b) By comparing Eq . (81a) to Eq. (82b) and Eq . (81b) to Eq . (82a) it follows that cdXE = Ady (83) The differentiation of Eq . (77) gives dX^ = -4n(Sz^F.do, so it can be deduced that , -4n6z. dy _ 1 _ do A (84)

PAGE 54

CHAPTER IV MATERIALS AND METHODS Goethite Preparation Three goethite preparations were made by mixing separately an equal volume (50 ml) of 1 N FeCCD^ with a similar volume of 2 N NaOH, 1 N NaOH, or 0.5 N NaOH in order to have OH/Fe mole ratios of 6, 3, and 1.5, respectively, in the mixtures. The suspensions were allowed to age for one week to induce crystal formation and growth. The addition of phosphorus to the goethite preparation was made for P/Fe = 3 ratio both at the beginning and by adjustment at the end of ageing process (one day of phosphorus reaction is allowed). In order to examine the effects of phosphorus concentrations on goethite crystallization, different levels of phosphorus in 2 N NaOH were mixed with 1 N Fe(Cl )2 a.nd allowed to age. After one week of ageing the suspension at either room temperature or at 55°C, the samples were washed by dialysis against distilled water during one week; the water was changed after intervals not greater than 12 hours. At the end of the ageing process and washing, the samples were freeze-dried . The existence of hydroxyl def ormational vibrations in the region 1200 to 1000 cm“^ was investigated by two methods i 41

PAGE 55

42 1) goethite and phosphated goethite were digested in D 2 O for 24 hours to replace the surface OH by OD , or 2) phosphated goethite samples were equilibrated with different salts (0.1 N KCl, 0.1 N KNO^, and 0.1 N NajSO^) and water to displace surface phosphates. The samples were separated from solution and dried for the infrared spectroscopy studies . An attempt was made to prepare FeOOD by digesting three times 0.9 g of FeC1^.6Fl2) in D 2 O and drying the suspension in order to make FeCl2.6D20. The residues were dissolved in 10 ml of D 2 O and mixed with 10 ml 2 M NaOD; the suspensions were then aged and dried as described above. Infrared and X-Ray Diffraction Techniques One milligram of the freeze-dried sample was mixed with 400 mg KBr to make pellets samples used for the infrared study. The infrared spectra are obtained by using PerkinElmer 567 grading infrared spectrophotometer. Prior to X-ray diffraction analysis using a general electric XRD700 instrument, and thin films of goethite samples were made on glass slides and allowed to air dry. Some Factors Affecting the Kinetics of Adsorption-Desorption The adsorption or desorption isotherm is obtained by plotting the amount of ion adsorbed or desorbed per unit v;eight of adsorbent versus the solution concentration. The

PAGE 56

43 solid and the solution were equilibrated through agitation for selected periods of time and temperatures. The equilibrium solution was removed after centrifugation and the phosphorus concentration determined through blue color development as ascorbic acid molybdophosphoric complex. Effects of pH The adsorption for different periods of times (0 to 72 hours) was done by equilibrating goethite or soils with phosphorus solution (Syg P/ml) or (20 yg P/ml) adjusted to pH 2, 4, 8, 9.5, and 11. Solid phase was separated by centrifuging at 5000 revolutions per minute (rpm) . Phosphate desorption was carried out on samples which reached the equilibrium state. The desorption was conducted in water, adjusted to the above pH, during 20 minutes, 6 and 12 hours. The pH of the solution was measured by placing the glass electrode in the supernatant after centrifugation. Effects of the Supporting Electrolytes The effects of the type of cations on the adsorption were studied by mixing 20 mg of goethite with 20 ml each of the phosphate solutions containing 1, 3, 5, and 10 yg P/ml in 0.01 M salts of NaCl, CaCl 2 , and AlCl^, each in separate experiments. With soils, the phosphate solution concentrations used were 0, 10, 20, 40, and 80 yg P/ml. The extent that each of the cations (Ca, Ha, and A1 ) contributed to the

PAGE 57

44 degree of reversibility of the adsorbed phosphate was studied by conducting the desorption with solution of the corresponding electrolyte (as in the adsorption) but without phosphate. A comparison of the phosphate desorption by different anions in various salts (KCl, K^SO^, KNO^, KCIO^, and NaOH) was made for goethite-P solution (one gram of solid with one liter of salt solution where the goethite had previously been treated with 5 ug P/ml). Time Effects on Adsorption and Desorption Adsorption of phosphates on goethite-solution containing 0, 1, 3, 5, 7, and 10 yg P/ml, and on soil-solution containing 0, 5, 10, 20, 40, and 80 yg P/ml was investigated after equilibration during different periods of time ranging from 1/10 hour to 80 hours. At the end of each reaction time, the solid-solution was centrifuged and the phosphate concentration remaining in the supernatant solution determined. The solid/solution ratio was 1/1000 for goethite-solution and 1/20 for soil-solution. Kinetic studies of phosphate adsorption were done at different conditions: 1) when different types of electrolytes were used, 2) for one electrolyte (CaCl 2 ) at different concentrations (.01 N Ca, 0.1 N Ca, 1.0 K Ca), and 3) at pH 2 and pH 10. After the adsorption proceeded until the equilibrium state v;as reached, desorption was conducted during 20 minutes, 6 and 12 hours in order to determine the minimum time required for complete

PAGE 58

45 desorption of this form of phosphate. This minimum time was also employed during desorption of phosphate when using either a desorbing solution at selected pH (2 to 11) or for different electrolytes. For those soils where it was observed that phosphate sorption did not follow the Langmuir isotherm, an attempt was made to determine the extent of changes in aluminum and iron phosphate by using the method of Peterson and Corey (1966). For aluminum phosphate, one gram of soil sample was washed with 2 N NaCl to remove exchangeable cations, the suspension was centrifuged and the supernatant solution discarded. Then 20 ml of 0.5 N NH^^F at pH8.2 were added and the suspension was shaken for an hour and then centrifuged for the P determination. For the iron phosphate fraction, the samples were washed with 2 N NaOH solution, centrifuged, and the supernatant retained for P determination . Surface Charge as Affected by Phosphate Adsorption The method of Lavardiere and Weaver (1977) was used to determine the net electric charge . The procedure was to add 20 ml each of 0.01 N CaCl 2 and 1.0 N CaCl 2 solutions per one gram of soil sample or 20 mg of goethite. Subsequent titrations of the suspensions were made with 0.01 N HCl or 0.01 N NaOH, by adding 0.1 to 0.3 ml at a time from a microburet at 2-minute interv^als. Continuous stirring was maintained and the pH read before the addition of either base or acid.

PAGE 59

46 A blank titration was made for the same volume of CaCl 2 solution. The amount of H or OH adsorbed at a given pH was calculated by the difference between the amount of H^ or OH added and that required to bring the blank solution of the same volume and salt concentration to the same pH as the soil or goethite. Soil pH, Iron Oxide, and Extractable P The soils used were obtained from Georgia (Cecil, Ap horizon), Colorado (roadside cut near Ft. Collins), and Kenya (latosol, sampled at 15-30 cm, near Kabete). The pH was measured in 1:1 soilto solution suspension for 10 g of soil with 10 ml of H^O or 10 ml of 1 N HCl , using a glass electrode. Iron oxide was determined by the dithionite-bicarbonate extraction and colorimetric determination of Fe as the ferrous orthophenonthroline complex. Extractable P was determined by three of the methods outlined by Ballard (1979). These data are shown in Table 1.

PAGE 60

47 Table 1. Soil pH, iron oxide and extractable P. Soils Georgia Colorado Kenya pHCH^O) 5 . 88 8.17 6.20 pH(KCl) 4 .95 7 .44 5 .75 % Fe202 2 . 24 0.67 3.95 P, ppm 0 . 05 N HCl +0.025 N 5.60 0 . 30 0.40 P, ppm 0.03 N NHj^F + 0.1_N HCl* 13.00 7.40 1.40 P, ppm 0.5 M NaHCOg-' 2 .80 1.20 1.60 Reagents described by Ballard (1974)

PAGE 61

CHAPTER V RESULTS AND DISCUSSION Goethite and Phosphated Goethite Studies by InfraRed Spectroscopy After the mixing of FeCI^ and NaOH solutions, some precipitates formed, but the particles formed at precipitation were not yet crystalline. The changes in precipitates to crystals originated from a discrete center or crystal nuclei (twinned and acicular crystals) produced by different mechanisms (Atkinson et al. 1968). The conditions governing the formation and nature of these crystal nuclei are not well known. Apparently, the number of crystal nuclei was increased by both temperature and hydroxyl ion concentrations. Iron and hydroxyl ions would be attracted to the centers of crystal growth as they lose their energies. As iron and hydroxyl ions were involved in the formation of the crystal, ageing reactions follow the sequence of lepidocrocite to goethite (Murphy et al. 1975). Another way would be that ferrhydrite forms first or just goethite was formed. The growth of the crystal occurred as the ions change phase during deposition from the solution. Hsu (1972) found that the removal of hydroxyl ions from the solution resulted in a drop in pH value of the solution as the time increased. 48

PAGE 62

49 Since the ions are changing phase from the solution (higher degree of randomness) to the solid phase (lower degree of randomness), the entropy change during the process from solution to crystalline phase must be negative. The entropy of ions within the crystal would be lower than those at the crystal surface. From infrared spectra, Kiselev and Lygin (1975) believed it was possible to estimate from adsorption entropy what was the degree of freedom of the adsorbed molecules . Goethite Structure Identification by Infrared The 4000 2000 cm ^ Region The goethite structure was interpreted through the identification of OH and FeO vibrations. The hydroxyl ion vibrations occurred at two strong bands in the 3700 2000 cm ^ region, centered at 3400 cm~^ and 3200 cm~^ (see Fig. 2 and Fig. 3). According to Nakamoto (1978), lattice water absorbed at 3550 3200 cm~^, he reported that the strong band centered at 3400 cm~^ was due to antisymmetric and symmetric OH-stretching of water. It could easily be recognized that the 3400 cm~^ band was not characteristic of goethite crystal structure since this band appears even when no goethite is present, as determined by X-ray diffraction which showed no crystalline material. It is of interest to observe that the intensity of the 3200 cm"^

PAGE 63

TRANSMITTANCE 50 3500 3000 2500 Frequency (cm Fig. 2. Infrared bands of (A) goethite ; (B) phosphate added to goethite at the end of ageing; and (C) phosphate added to goethite at the beginning of ageing. Curves A, B and C are for 22° C and curves A', B', and C are for 55° C.

PAGE 64

TRANSMITTANCE 51 Infrared bands of (A) goethite ; (B) phosphate added to goethite at end of ageing; (C) phosphate added to goethite at beginning ageing. Curves A, B, and C are for 22° C, and curves A’, B', and C are for 55°C. Fig. 3.

PAGE 65

52 band decreased as the OH/Fe ratio in the ageing solution decreased. The band centered at 3200 cm”^ appeared only if goethite structure was present. As a result, it was concluded that the 3200 cm ^ band corresponded to structural Fe OH stretching. The 2000 3000 cm ^ Region After goethite was freeze-dried, not all the adsorbed water was removed. Nakamoto (1978) indicated that the relative velocity of the oxygen nucleus compared to that of hydrogen nucleus is small. This meant that the surface binding of water through the oxygen atom to goethite would not induce a significant change in the overall water vibraTion which could be observed at around 1620 cm ^ as depicted for water structure below /k H H d N /\ H H H H Case V. Case V, Case V. The above three normal modes of vibration in H^O are infrared active. The bending vibration ^ is centered at 1620 cm ^ and the water stretching bands (v^ and v^) vibrate in the 3400 cm region .

PAGE 66

53 The 1300 700 cm ^ Region The two main bands at 890 and 790 cm"^ would be assigned to Fe-OH bending vibration of the structural hydroxyls. Busca et al. (1978) supported this finding by observing that the structural OH in a-FeOOH disappeared upon heating to form a-Fe^Og. The appearance of the 890 and 790 cm~^ bands always indicated goethite crystallization, which can be observed in Fig. 3, 4, 5, 6, 7, 9, 10 and 11. From the above mentioned figures, it was observed that for each OH/Fe ratio, the way in which phosphate was added had an effect on the degree of goethite crystallization. From Fig. 5, as reported by Farmer and Palmieri (1975), typical goethite could be identified by the Fe OH bending vibration at 890 cm~^ and 790 cm"^. In Fig. 3, where OH/Fe = 6, there are also strong bands at 890 and 790 cm~^ in all cases, except when phosphate is added at the beginning of ageing of the suspension at 55°C. In the latter case, the increase of temperature to 55°C favored preferential phosphate binding to the iron which was observed by the strong vibration at 1000 cm As shown in Fig. 3c, phosphate was within the goethite structure because there was both the P 0(Fe) vibration at 1000 cm ^ and the surface binuclear (FeO^POOH vibrations at 1190, 1100, and 1030 cm“^. When CH/Fe = 3, it was not possible to have phosphate v;ithin the crystal because there was no crystallization at room temperature. However

PAGE 67

TRANSHITTAKCE 54 Frequency (cm Fig. 4. Infrared bands of (A) goethite ; (B) phosphate added to goethite at end ageing; and (C) phosphate added to goethite at beginning ageing. Curves A, B, and C are suspensions at 22°C, and curves A', B', and C are for suspensions at 55°C.

PAGE 68

TRANSMITTANCE 55 Fig. 5. After_ Farmer and_Palmieri (1975). Infrared bands of Goethite and lepidocrocite .

PAGE 69

TRANSMITTANCE 56 Fig. 6. Infrared bands of (A) goethite; (B) phosphate added to goethite at end ageing; and (C) phosphate added to goethite at beginning ageing. Curves A, B and C are for suspensions at 22°C and curves A', B', and C are for 55°C.

PAGE 70

TRANSMITTANCE 57 Fig. 7,. Infrared bands of (A) Fe hydroxide material; (3) phosphate added to Fe hydroxide material at end of ageing; and (C) phosphate added' to Fe hydroxide material at beginning ageing. Curves A, B, C are for suspensions at 22° C and curves A’, B’ , C are for suspensions at 55° C.

PAGE 71

58 there was weak crystallization at 55°C that favored some phosphate surface binding (see Fig. 6b). In OH/Fe = 1.5 suspension, after one week of ageing either at room temperature or at 55°C, no goethite formation was observed. There was a single, strong vibration at 1030 cm ^ when phosphate was added at the end of the ageing, but the vibration was at 1000 cm ^ when the phosphate was added at the beginning of ageing. The 1030 cm~^ vibration, which is the P OH bending vibration, indicated the excistence of phosphate at the surface, and the presence of the 1000 cm~^ vibration from the P 0(Fe) showed phosphate directly bound to iron when phosphate was added just prior to the ageing process (Parfitt et al. 1975). Goethite Identification by X-Ray The X-ray diffraction peaks for goethite were at 4.19 8, 2.70 8, and 2.45 In all cases, the 4.19 X is the most intense peak while those at the 2.70 ^ and 2.45 K were weak. Fig. 8. The presence of the 2.70 8 spacing indicated that some hematite might be present. However, because the 2.45 8 peak was present, it was believed that considerable amounts of goethite existed (Schwertman and Taylor 1977). According to their studies, it would be possible for hematite to exist along with goethite because the hydration of hematite would yield goethite with a standard free energy (AG°) of the reaction varying from -0.2 to 0.4 kcal/mole.

PAGE 72

Angstrom Spacing (8) 59 01 4 -> • •H bO rC: c P -H 0 ) 0 ) O M hO fO Degree (20) Fig. 8. X-ray diffraction patterns of (A) goethite, (B) phosphated end of ageing, and (C) phosphated goethite at beginning of

PAGE 73

60 Factors Affecting Goethite Crystallization Effect of OH/Fe Ratio The above results indicated that the bands at 3400 cm” 3200 cm 890 cm and 790 cm ^ were characteristic of the appearance of goethite. For the same ageing period, an increase in OH/Fe favored goethite crystallization. When the OH/Fe is 6, the bands at 3200 cm”^, 890 cm”^, and 790 cm were strong, indicating that goethite structure was wellformed. Where OH/Fe is 3, after a week of ageing, goethite structure was apparently present only if the suspension was kept at 55°C, even then the degree of crystallization was less intense at OH/Fe = 3 than it was where OH/Fe = 6. This weak crystallization was suggested by the weak band at 3200 cm ^ (which took the form of a shoulder) and by the less pronounced intensities of the 890 cm ^ and 790 cm ^ bands. However, when OH/Fe is 1.5, there was no goethite crystallization even when the suspension was aged at 55°C. The above conclusion disagreed with that of Atkinson et al. (1974) who claimed made goethite was in a suspension at 28°C where OH/Fe is 2.0 when aged for 50 hours. Fhosphated goethite The application of phosphorus weakened the goethite structure because some phosphate was probably incorporated within the lattice. In Fig. 4 and Fig. 6 where OH/Fe is 3,

PAGE 74

61 the addition of phosphate to goethite suspension resulted in bands at 3200 cm 890 cm ^ and 790 cm ^ which were weaker than those for OH/Fe = 6 after ageing. When OH/Fe = 6 in the suspension, hydroxyl ion concentration was high enough to form the necessary bonds for the goethite structure to appear. The presence of phosphate on the goethite surface could be recognized by the appearance of bands in the 1200 1000 cm ^ region, but vibrational hydroxyl deformations also could occur in this same region. Parfitt (1979) stated that the P = 0 bond had stretching vibrations in the 1190 cm ^ and the 1030 cm ^ region. When the phosphate was added at the end of the goethite ageing, there was an appearance of P = 0 vibration at 1190 cm ^ and the 1030 cm ^ due to P OH vibration in agreement with work by Parfitt (1979). At high temperature (55°) even when OH/Fe = 6, the phosphate is preferentially bonded to the iron, thereby reducing the capacity of hydroxyls ions to bind freely with iron to form goethite. As a result, there was direct binding of phosphate to iron which was confirmed by the strong band at 1000 cm~^ assigned to P 0(Fe) and by the lack of the Fe OH bending vibrations at 890 cm”^ and 790 cm~^. This is confirmed by spectra given in Fig. 3. In Fig 9, Fig 10 and Fig 11 where suspension the OH/Fe ratio was either 6, 3, or 1.5, the increase in phosphate concentrations relative to the iron (P/Fe = 0.032,

PAGE 75

TRANSMITTANCE 62 i.'ave Number (.cm “ ) Fig. 9. Infrared bands of phosphated goethite at beginning of ageing for suspensions of various P/Fe values at OH/Fe = 6.

PAGE 76

TRANSMITTANCE 1200 1000 800 600 400 V.'ave Number (cm . Infrared bands of phosphated goethite at beginning of ageing for various P/Fe ratios at OH/Fe = 3.0. . 10

PAGE 77

TRANSHITTAIICE £4 1200 1000 800 600 400 VJave Number (cm~^) Fig. 11. Infrared bands of phosphated goethite at beginning of ageing for various F/Fe ratios at OH/Fe = 1.5.

PAGE 78

65 0.32, 3.2) induced an increase in the intensity of P 0(Fe) and P OH stretching vibration at 1000 cm~^ and 1030 cm~^, respectively. When P/Fe was greater or equal to 0.3, these two bands overlapped, showing only one very strong band at 1000 cm~^. In order to identify the vibration of hydroxyl deformation assumed to be in the 1200 1000 cm"^ region, two methods were used: 1) goethite and phosphated goethite were digested in D 2 O for 24 hours and 2) surface phosphate was desorbed by different anions and water. After the samples were dried, the infrared spectra were similar to those shown in Fig. 12 and Fig. 13. The weak bands due to OH deformation were displaced by OD and only the P 0(Fe) stretching vibration at 1000 cm~^ and those for the Fe OH bending at 890 cm ^ and 790 cm ^ persisted. Evidently, it could be assumed that the bands at 890 cm”^ and 790 cm~^ were due to structural Fe OH bending which cannot be affected by digestion in as long as the initial goethite maintained its structure. In a further study, the evidence for hydroxyl deformations was examined when the phosphated goethites were desorbed by different salts (KCl, KNO^, Na2S0^^) and water. After either 0.1 N KCl solution or water desorption, hydroxyl deformations were not removed in the 1200 cm ^ 1000 cm ^ region. With 0.1 N Na2S0^ solution used for desorption, the infrared spectra showed strong

PAGE 79

TRANSMITTANCE 66 Fig. 12. Infrared bands of 1) goethite digested in D„0 and 2) phosphated goethite digested in D^O.

PAGE 80

TRANSMITTANCE 67 '.v’ave Number (cm ) Fig. 13. Infrared bands of phosphated goethite after desorption by (1) 0.1 N KCl, (2) H„0, (3) 0.1 N KNO,, (4 ) 0.1 N Na^SO^^.

PAGE 81

68 2 -1 bonds of SO^ ions in the 1200 1000 cm region. When 0.1 N KNOg was used for desorption, NO” appeared to displace some surface phosphates so that the hydroxyl vibration in that region was reduced (Fig. 13). It can be said that NO^ 2 . and ions have the ability to desorp readily displaceable phosphate at the surface. However, sulfate ions could not be used to provide evidence for this type of phosphate desorption by infrared since sulfate ions have strong vibration in the same region as that of surface-bound phosphate. Parfitt et al . (1975) assigned the appearance of weak bands in the 1200 1000 cm ^ region for non-phosphated goethite to the Fe OH deformat ional vibration which can partially overlap the region of P = 0 stretching and P OH bending vibration. The question arose about the type of hydroxyl groups that are displaced when phosphates are added to the goethite suspension. Russel et al . (1974) indicated that there were three types of hydroxyl groups: 1) where OH is singly coordinated to an Fe atom with hydrogen bond interaction with another atom, 2 ) where OH is coordinated to tv/o Fe atoms and 3) when OH is singly coordinated to an' Fe atom. Parfitt et al . (1975) considered that only type 3 of OH was displaced by phosphate while OH of the type 1 and 2 are unreactive. The types 1 and 3 of OH, they considered as the structural OH v;ith three vibrations at 3200, 890, and 790 cm

PAGE 82

69 type 3 of OH Fe -f OH Fe OH Fe I OH [001] face Fig. 14. The 001 face of goethite lattice. (After Bragg and Claringbul 1965). The formation of binuclear bridging resulted from the displacement of two adjacent type 3 hydroxyls Fe 1 Fe Fe Fe 1 1 OH OH OH 1 OH \ / \ / ^Fe "Fe 1 1 \ / Fe 1 0 0 OH Fig. 15. The 001 face of phosphated goethite. Phosphate Adsorption and Fesorption Studies Effects of the Initial Concentration The phosphate adsorption isotherm is obtained by plotting the amount of phosphorus adsorbed (yg P/g of adsorbent)

PAGE 83

70 against the equilibrium concentration of phosphorus in solution. The shape of each isotherm curve is altered by the relative amount of phosphorus adsorbed at each equilibrium concentration and reaction time . For a particular time of adsorption reaction, the amount of phosphorus adsorbed increases with the initial concentration but the percentage of phosphate sorbed decreases. For the goethite-P solution system, the shapes of the phosphate adsorption isotherm curves were affected by the relative amount of phosphorus adsorbed at each equilibrium concentration (Fig. 16). Each curve was composed of three parts: one at low solute concentration (< 0.25 yg P/ml), a second part at higher solute concentration (0.25 to 1.5 yg P/ml) where the isotherm becomes convex, and a third linear part at higher solute concentration (> 1.5 yg P/m.l). However, in the soil-solution system, phosphate adsorptive capacity depended on soil characteristics which affected the shape of the isotherm curves. The Kenya soil and Georgia soil phosphorus sorption curves apparently had two portions, the first part which is a curved portion (< 5 yg P/ml) and a second or linear portion (> 5 yg P/ml). In Fig. 17, the apparent lack of an initial linear portion for the curve found for Kenya soil suggests need for more adsorption data at low equilibrium concentration (< 3 yg P/ml). For the Georgia soil, (Fig. 18), the adsorption maximum is low compared to

PAGE 84

71 E O. bO C O 4-> C ( 1 ) a c o CJ E •rH XI •H 3 a" U5 C o •H P I — I 0 to 1 QJ P •H X P 0) o bO P O P t/3 QJ E *H P c O *H P O 0) > X) (U p o 0) p p rti W ftJ U] e 0) P o w c o •H • P W O. E QJ O P w w XJ >, < w (a:iTmao 3 3/j 3ui) uoiadjos bO *r-| P

PAGE 85

Sorption (100 pg p/g Soil) 72 Fig. 17. Adsorption isotherm for the Kenya soil after 24 hours of reaction time.

PAGE 86

73 e bO 3. c o •H P P C 0) o c o CJ e zi •H •<— I »p rH o cr u (TT°S 3/d 3ri OOT) uoTa.daos Fig. 18. Adsorption isotherm for the Georgia soil after a reaction time of 24 hours.

PAGE 87

74 c o •H •P X-J C Q) U C O O e zi XI •r*H o" UJ O w o XI fO o I — I o u QJ x: M o IM e u ( 1 ) x; p o in C o H p • Cl, 0) p e O -H w p> XI iti c o 0) -H X P> f 4 O X t 4 P, (U W p O X X iP o bO •H

PAGE 88

75 that for the Kenya soil, which means that for the same initial concentration in solution more phosphate ions are present in the Georgia soil than in the Kenya soil. For the Colorado soil, (Fig. 19), there is a slight change in slope for the initial and final linear party of the absorption isotherm. For the initial range of equilibrium concentration ranging from 0 to 25 yg P/ml, the amount adsorbed increased linearly as the equilibrium concentration increases. This relationship was 20 yg P adsorped per gram of soil for each yg P/ml. Muljadi et al . (1966) noted that their isotherm curves were also linear initially with a curved transition to the second linear portion. They ascribed the initial and curved portions to phosphate exchange with OH of Al(OH) located on the clay edge surfaces. They do not give a clear explanation of the mechanism of adsorption reaction for the third part of the curve but postulated that the final linearity of the isotherm indicated that the number of adsorption sites remained constant even though the amount of phosphate adsorbed increased. The time required for the reaction to reach equilibrium decreased as the initial concentration decreased, (Fig. 20). If the initial concentration was less than or equal to 5 yg P/ml, the steady state of reaction was reached in less than 10 hours. Tv;enty hours of reaction were required to approach the equilibrium state with an initial concentration

PAGE 89

10 MR p/ml 76 (9q.Tq3.a03 §/j auj) uoxqdaos Fig. 20. Effects of the initial P concentration on the phosphate adsorption by goethite-solution (1 g/1000 ml).

PAGE 90

77 greater than or equal to 5 yg P/ml. The above observations showed that low amounts of phosphorus were almost instantaneously adsorbed onto the synthetic goethite. Because the initial phosphate potential on the solid phase was low or nil relative to the phosphate potential in solution, the phosphate flux from the solution was therefore high. This movement of phosphate to the solid surface continued for a longer period of time if the initial concentration (or phosphate potential) in solution was high. Effects of the Supporting Electrolyte Type of Electrolytes The adsorption isotherm of phosphate on geothite was examined using molar concentration of the electrolytes, NaCl, CaCl^ , and AlCl^ , as given in Fig. 21. Where the initial concentration was greater than 0.5 yg P/ml, the salts (electrolyte) gave a significant effect on the amount of phosphorus adsorbed. The adsorption at each concentration decreased in the order IM A1C12>1M CaCl 2 >lM NaCl. It was noted that the adsorption of phosphate using water as supporting medium gave the same adsorption isotherm as that using 1 M NaCl solution. For equilibrium concentration greater than 1 ug P/ml, the amount of P sorbed in 1 K AlCl and 0 1 M CaCl^ was greater than that sorbed in 1 H NaCl by a factor of 1.6 and 1.5, respectively.

PAGE 91

H AICI3 78 (9a.7qq.e01j 3 /d Sm) uoiq.daos Equilibrium Concentration (qg P/rrd) Fig. 21. Effects of the type of supporting electrolyte on P adsorption on goethite-suspension (1 g/1000 ml).

PAGE 92

79 Table 2. Type of salt effects on the P adsorption on goethite at 5 pg P/ml. Electrolyte P adsorption (pg P/g) 1 M NaCl 3780 1 M -CaCl^ 5600 1 M A1C1„ 5900 — 3 Ryden and Syers (1975) reported that, where some soils have a final concentration above 0.1 pg P/ml, the P sorption _ 2 in 10 M Ca solution was 1.5 to 2.5 greater than the sorption from water. From the isotherms for P sorption by goethite, the data were arranged as the plots of a/b versus a, where a is the equilibrium concentration, b is the amount of phosphorus adsorbed per unit weight of adsorbent. The linearity of the plots (Fig. 24) confirmed the Langmuir type of adsorbent. When using Eq . (14c) the adsorption maximum (s) and the adsorption energy constant (k) are calculated as shown in Table 3.

PAGE 93

80 Equilibrium Concentration (pg P/ml) Fig. 24. The transformed Langmuir equations for phosphate sorption by goethite as affected by reaction times, a is equilibrium concentration and b amount adsorbed.

PAGE 94

81 Table 3. Effects of three electrolyte salts on the phosphate sorption maximum and sorption energy constant for goethite. Electrolyte Adsorption Maximum Sorption Energy Constant (yg P/g) (ml/yg P) 1 M NaCl 4300 3 . 8 1 M CaCl^ 6900 2.3 1 M AlCl^ 12700 1.1 The increase of phosphate sorption maximum, (Table 3), was accompanied by a decrease in the apparent sorption energy constant. The increase in phosphate sorption may be due the increase in the cation charge of the electrolyte favored adsorption. A probable way in which the cation charge enters into the phosphate adsorption reaction was through cation bridging: Fe 0 so that where M is Na, then n = 0 , or if M is Ca then n = 1, and if M is A1 then n = 2. As the charge on the cation increased, there v:as a greater attraction between the cation

PAGE 95

82 on the goethite surface and the phosphate ion. The work required to bring the phosphate ion to the cation evidently decreased as the valence increased. The increase in metal valence favored a higher probability that the phosphate was maintained in the goethite-P solution interface, so that the energy for adsorption decreases. Both aluminum and calcium ions also probably reacted with phosphate in solution to form new phases (precipitation), inducing thereby a decrease of phosphate in solution. Som.e of the compounds which might precipitate in solution depended both on phosphate concentration (molarity) and pH. These systems can be written as follows: For CaHPO^ , the pH and H2P0^ relation is pH2P0^^ = pH 3.14 (86) For A1 ( OH ) ^H^PO^ , the pH and H^PO^ relation is pH^FO^^ = pH + 10.7 (87) For Fe (OH ) ^H^PO^ , the pH and H^FO^^ relation is pH^PO^^ = pH 10.9 (88) The plots of pH^PO^^ versus pH gave the solubility diagrams of the compounds as illustrated by Lindsay and Moreno (1960). At any point (pH, pH^FO^) above the line for a selected compound, precipitation is expected while below the lines dissolution of the corresponding solid phase will occur. In our study, the pH ranged from 6 to 7 so that with

PAGE 96

83 a phosphate concentration of 5 yg P/ml or 1.72x10“^ moles of H^PO^ per 1 liter, we would have pH 2 P 0 ^^ = -log H^PO^^ = 3.8; hence, no important amount of precipitation of the above solid phase was expected. Other complex compounds involving combination of Ca and Fe phosphates or A1 and Fe phosphates could precipitate near or on colloid surfaces of the goethite or the soils. Effects of Electrolyte Concentration Salt concentration had a marked effect on the amount of phosphate sorbed during reactions at various periods. The time at which equilibrium was reached was evidently not affected by increasing the salt concentration (Fig. 22). When the initial concentration is 5 yg P/ml, at equilibrium state the phosphate sorption in 1 N CaCl 2 and 0.01 N CaCl 2 is 1.3 and 1.2, respectively, greater than that found for water system without salt. Van Olphen (1977) pointed out that increasing the electrolyte concentration not only caused compression of the diffuse part of the double layer but also some ions, as counter ions, shift from the diffuse layer to the Stern layer. As a result, the ion concentration of the diffuse layer decreased and more adsorption took place. This concept was supported by Ryden and Syers (1975) who found that equilibrium phosphate concentration remained 2 -1 the same in both 10 M Ca and 10 M Na systems. Donnan theory could be used to explain the above phenomenon. The inner solution possibly ranged from the solid surface to the

PAGE 97

84 Fig. 22. Effect of the supporting electrolyte concentration on the kinetics of F adsorption by the goethitesolution system (1 g/1000 ml).

PAGE 98

85 Table 4. Phosphate desorption from goethite by different anions . Anions (O.IN) Cl 2 SO 4 no; cio; OH H^O % Adsorbed P desorbed 0 . 3 1.4 1 — 1 1 — 1 CO 1 — t cn CO 2.5 Table 5. Effects of reaction time on phosphate adsorption maximum and sorption energy constant for the goethite system. Reaction time ( hours ) Adsorption Maximum (ug P/g) Sorption Energy Constant (ml/yg P) 1/10 1340 0.6 6 1/2 2280 0.84 8 3850 2.40 18 to 76 4460 4.60 Table 6. The logarithm of (a) equilibrium P concentration as a function of the logarithm of (b) amount of P adsorbed. log a -.45 .15 3.07 5.54 log b 3 . 40 3.55 3.60 3.65

PAGE 99

86 imaginary limit of the physically adsorbed ions and the outer solution containing all free ions. This could be expressed as: = (M)J/''(H2P0^^). = k (89) where (M)^ and (M)^ are the concentrations of the cation M of valence n in the outer solution and inner solution respectively, terms (H^PO^^)^ and (H 2 P 0 ^)^ are the phosphate concentrations in the outer and inner solutions, respectively, and k is a constant. From the above equation, it was evident that as salt concentration (M)^ increased, the (H 2 pO^^)^ concentration must decrease. As a result, phosphate may be adsorbed through cation bridging or precipitation as cationphosphates. Phosphate Desorption Effects of phosphate free solutions of 1 M AlCl^, and CaCl 2 and NaCl are shown in Fig. 23. There was a significant effect of the type of salt on the extent of phosphate desorption. The magnitude of differences due to salt sources increased as the equilibrium phosphate concentration increased. For a low equilibrium concentration (1 yg P/ml), the desorption with 1 M AlCl^ solution released 550 yg P/g of goethite. This was 3.1 and 1.8 times greater than similar desorption by 1 K CaCl 2 and 1 K MaCl, respectively. The higher desorption of phosphate by 1 M AlCl^ solution agreed with the apparent lower potential binding energy constant of

PAGE 100

Desorption ( pg p/g Goethite) 87 rquilibrium Concentrntion (Ug P/ml) Fig. 23. Phosphate desorption in 10 hours from goethite by three salt solutions at various equilibrium F concentration.

PAGE 101

88 phosphate in goethite-P solution in the presence of this salt. The increase in phosphate adsorption due to increase of cation valence was also accompanied by an increase in the desorbed fraction which was weakly bound (Fig. 23). This would suggest that the additional phosphate adsorption in presence of multivalent cations may not only be due to direct surface binding but also to other bonding by some form of cation bridging and/or multilayer phosphate adsorption. These latter forms were easily replaced during P desorption with further fresh salt solutions. Time of Reaction Effects Adsorption Pxeaction From the decrease in the initial phosphate concentration, the phosphate sorption increased with time and remained constant when the reaction time was greater than 18 hours for goethite-P solution system (Fig. 25). The adsorption made according to the transformed Langmuir equation, Eq. (14c), was linear. This suggest that the surface of the synthetic goethite was composed of an homogeneous population of sites for the adsorption of phosphate. In Fig. 24, it can be seen that both the slope and intercept of the sorption isotherms decreased as the reaction time increased up to 18 hours. In Table 5 it v;as shown that both the calculated adsorption miaximum and the potential binding energy constant

PAGE 102

89 w p o x:
PAGE 103

90 (-[u/j 3rt) uoTitaiusDUO^ 'jjnTaqi -[TnbH Fig. 26. Change in the equilibrium P concentration for three soils as affected by the reaction times.

PAGE 104

91 increased as the reaction time increased. Changes in this potential binding energy constant increased as the reaction time increased. Changes in this potential binding energy might be explained by the fact that as the surface becomes covered by phosphate, further adsorption of other phosphate ions occurred on the remaining free surface which required mueh more energy for bonding than on the first monolayer (Clark 1970). It should be kept in mind that the Langmuir equation assumed that there was only a monolayer of adsorbed P and that there was no lateral interaction between the adsorbed ions. As a result, the application of the Langmuir equation to experimental data would require that the reaction time be specified since the nature of the adsorbed layer would not be defined with time. The increase of energy constant as the adsorption maximum increases was not in disagreement with studies by Aharoni and Ungarish (1977) who stated that adsorption took place in the initial stages on the sites with lowest energy at a particular time. When comparing materials of different mineralogical composition, Ryden and Syers (1975) found that consistent trend for increase in the energy constant with the adsorption maximum confirmed the assumption that phosphate was usually adsorbed onto all available low energy sites before the high energy sites were filled. In the present studies, using the P sorption data for goethite-P

PAGE 105

92 solution when the reaction times exceeded 18 hours, further treatment of the data was made using the transformed Freundlich relationship given in Eq . (10). From Table 7, the plot of log a versus log b was close to linearity as the following regression equation: log b = 0.209 log a + 3.50 (90) Therefore, the free energy of adsorption (expressed as the energy difference between the adsorbed and free state of phosphate) decreased exponentially as the binding energy and surface coverage increased. From the proposed kinetic model given in Eq. (57) and Eq . (59b), an attempt was made to predict the change with respect to time that occurred for forms of phosphate logically in the system. These were (A) free ions such as H^PO^, (B) the physically and reversible chemically adsorbed phosphate, and (C) the irreversible chemically adsorbed phosphate. A transformation of Eq . (57a) can be written as In a = w^(t t^) + k (91) where a is the equilibrium concentration, w^ and k are constant, t is the reaction time, and t^ is the equilibrium time equal to 18 hours. According to the sorption, data were plotted using Eq . (91) and are given in Fig. 30. It can be seen that the phosphate adsorption reaction on gpethite followed two steps before the equilibrium was reached. The first step was finished after 3 hours of reaction tim.e (t^).

PAGE 106

In (Equilibrium Concentration [pg P/ml]) 93 ' Fig. 30. Logarithmic plot of equilibrium P concentration change with time (t) relative to equilibrium at 18 hours (t^) for phosphated goethite. Data were obtained using Eq . (91).

PAGE 107

94 and the second step after 3 to 18 hours (t^). From these data it was apparent that Eq. (57a) should be split into two components with respect to time after initiating the experiment so that a, = a e 1 o (92a) and for the range t^ to t^ a , e ol w,, (t t ) lb o (92b) so that a^ is the initial P concentration in solution for t ^ and a 2 is the solution P concentration at longer reaction time between t^^ and t^. In the goethite-P solution system, there was an initial fast reaction so that w la = hour ^ and a second step of the reaction which was slower with = 0.025 hour This would also apply for the fast and slow adsorption reactions observed by Barrow and Shaw (1975). Desorption of Phosphate After the phosphate adsorption was conducted for different periods of time, the desorption curve shown in Fig. 27 for the Kenya soil indicated that the phosphate desorption in 0.01 N CaCl^ was almost linear with respect to P concentration in the equilibrium solution and that the adsorption reaction tim.e affected the desorption process. With an initial concentration of 20 ug P/ml after 23 hours

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Fig. 27. Effect of reaction time and equilibrium F concentration on P desorbed by 0.01 b CaCl^ from Kenya soil.

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96 of adsorption reactions in 0.01 N CaCl 2 solution desorption was 80 ug P/g for Kenya soil. Similarly at each equilibrium concentration, the desorption of phosphate increased for the Georgia soil. Fig. 28 also showed more desorption of P from the Colorado soil after 23 hours of P adsorption than that after 27 hours. For the Colorado soil, the phosphate desorption data indicate that if the equilibrium concentration was greater than 5 pg P/ml, there was a divergence in phosphate release as the adsorption reaction time increased. In Fig. 29 extraction with 0.5 N NH^^F at pH3 . 8 resulted in a significant increase in phosphate release, presumably from that in aluminum form. Apparently, more phosphate was held as an aluminum phosphate after 72 hours of adsorption than after 21 hours, a negligible amount being recovered after 4 hours of adsorption time. This might explain the lower desorption of phosphate by 0.01 N CaCl 2 at longer adsorption time . In the goethite-P solution system, (Fig. 25), phosphate adsorption at first increased and then remained constant as the reaction time was increased. This indicated that there is first a predominant physically adsorbed P followed by a conversion to chemically adsorbed P; this is supported by the increase in P sorption energy constant (Table 5). The free phosphate will replace some of the physically adsorbed phosphate until a stable layer will be established and the system is then at equilibrium.

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Desorption (pg p/g Soil) 97 Fig. 28, Effect of reaction time and equilibriumi F concentration on P desorped by 0.01 N CaCl„ from the Georgia and Colorado soils.

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98 Fig. 29. Effect on P desorbed by 0.5 N NH^F as affected by the adsorption reaction time and equilibrium P concentration .

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99 Effects of pH on P Adsorption and Desorption Adsorption The adsorption of phosphates both on goethite and soils induced an increase in the equilibrium solution pH, (Table 7). The calculated pH change was 0.75xl0”^pH units/pg P/g of goethite in the soil-P solution systems; it was 3.75X10" 2.14X10 and 0.75X10 ^pH units per pg P/g soil for the Georgia, Kenya, and Colorado soils, respectively. In the goethite-P solution studies, phosphate adsorption increased with decreasing pH of the equilibrium solution (Fig. 31). When the initial P solution was 5 pg P/ml, the amount of phosphate adsorbed decreased almost linearly with pH of the equilibrium solution was increased according to the equation (pg P/g) = -446 pH + 5680 (93) where P is the amount of P absorbed on goethite. The effects of pH on the prevalence of the orthophosphate species in solution have been investigated. From Lindsay and Vlek (1977), each of the following phosphate species constituted more than 50% of the total P at the following respective pH range: i) H^PO^ for pH 2 or less, ii) H^PO^^ for pH 2 to 7 , iii) HPO^ for pH 7 to 12, and 3 iv) P0[^ for pH greater than 12. Such changes in phosphate ions with pH further complicated the interpretation of phosphate adsorption. As the pH increased, the goethite had an increasing net negative charge (Fq. 64). Since the

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100 Fig. 31. Effect of change in equilibrium solution pH on P adsorption and 0.01 N CaCl_ desorption of P in goethite system. Initial P concentration is 5 yg/m.l.

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101 Table 7. Effects of solution pH P ' adsorption on after 24 hours the equilibrium of adsorption reaction Adsorbent P added (pg P/g) P adsorbed (yg P/g) pH Initial Final Goethite 5 3500 7.10 7 .30 10 4300 6 .60 7 . 00 Georgia 5 179 5.10 5 .55 10 202 5.10 5.47 Colorado 5 198 6 .80 7 .03 10 374 6 .57 6 . 69 Kenya 5 304 5.50 6 . 06 10 490 5 .40 5.95

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102 predominant orthophosphate ions are in the increasing order 23HPO^ , PO^^ with increasing pH, from the electrostatic point of view, there was an increasing repulsion between the goethite surface and the orthophosphate ions. This resulted in less phosphate adsorbed with an increase in pH. Hingston et al. (1967) found that the maximum adsorption of goethite was approached when the pH was less than 4. At such low pH, goethite would have a net positive surface and phosphate ion would be negatively charged. As a result, minimum repulsion or maximum attraction between goethite surface and phosphate ions should occur at this acidity. Desorption The reversibility of the adsorbed phosphates from solution at different pH condition is shown in Fig. 31. The pH at which the phosphate adsorption took place had an important effect on the reversibility of phosphate adsorption. As the pH of the equilibrium solution (adsorption process) increased, there was first a decrease in the quantity of phosphate desorbed until a minimum was reached in the range pH 5 to pH 6 followed by descreased desorption as pH was further increased. At less than pH 5, both phosphate adsorption and desorption were greately increased. This indicated that, as the pH decreased, there were some additional positive sites formed on the goethite surface. One probability was the shift of hydronium ion to the hydroxyl group of Fe OH as shov;n below

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103 OH H ' Fe OjJ+ . . . 0 P = 0 (94) OH resulting in more phosphate adsorption. However, when the pH was greater than 6, there was an increase in the hydroxyl ion concentration which has strong affinity for any positive site. Consequently, the existence of competition between OH and the orthophosphate ions for the same type of sites resulted in lower phosphate adsorption. Hingston et al. (1968) noticed that anions were desorbed by competitive effect when the competitor can occupy sites in addition to those already occupied by other anions, hence the increase in the negative charge of the surface. This increase in negative charge produced by the competitor allowed phosphate desorption by surface hydroxyls. It was known that the increase in OH concentration favored the formation of surface negative sites, which contribute to the weak phosphate binding resulting in greater P desorption. The competitive capacity of hydroxyl for phosphate sites was further confirmed from data shown in Fig. 32. The pH effect on the equilibrium P concentration after various adsorption times was m.arkedly different for pH2 and pHlO. At pH2, the desorption of phosphate increased initially and remained constant after 5 hours of reation tim.e. However at pHlO the phosphate desorption increased sharply initially

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Desorption (ug P/g Goethite) 10 ^ Fig. 32. Effect of phosphate adsorption time on phosphate desorption by 0.01 N CaCl 2 at pH 2 and pH 10.

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105 and thereafter decreased slowly and linearily as the reaction time increases. This shows that at pfIlO the hydroxyls ions continued the phosphate desorption as the adsorption reaction time was increased. This apparent increase in efficiency for OH replacement of phosphate might also have included calcium phosphate precipitation as well as the fact that much of the phosphate ions were probably in trivalent state. Obviously, ageing of the phosphated goethite negated phosphate desorption by OH replacement. Effects of pH and Time In order to define the response of phosphates desorption with time and pH, the desorption was conducted at both different periods of time and pH. From the data shown in Fig. 33, at both 6 and 12 hours of desorption time, phosphate desorption followed the same pattern over the same pH range. However, after 20 minutes, only at pH 10 was substantial amounts of phosphate desorbed (600 yg P/g goethite) which is about 1/3 of the amount that was obtained after 12 hours of desorption at this pH. Generally, it can be said that the efficiency of phosphate desorption increased as the pH of the desorbing solution was increased. The effects of the electrolyte concentration and pH interaction on phosphate desorption are shown in Fig. 34. The phosphate desorption at any pH increased when supporting electrolyte was present compared to use of water only for

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Desorption (mg P/G Goethite) 106 Solution pH Fig. 33. Effects of solution pH on the amount of F desorped from goethite. t>

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Desorption (mg P/g Goethite) 107 Fig. 34. Effects of pH and concentration of supporting electrolyte on phosphate desorption.

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108 phosphate desorption. In the range pH 2 to pH 12, either 1 N CaCl 2 or 0.01 N CaCl 2 gave similar amounts of desorption phosphate at each pH, and each is greater than that obtained from water. At any pH, the phosphate desorption in 0.01 N CaCl 2 was approximately 200 yg P/g goethite greater than the amount desorbed in water. Effects of P Adsorption on Surface Charge of Goethite In goethite-solution and soil solution system, it was assumed that H and OH ions are potential determining ions. Potentiometric titration of goethite was made to determine the zero point of charge (ZPC) from the distribution of the net electric charges with varying pH and electrolyte concentrations (Fig. 35). The ZPC was at the inflexion point of the titration curves at different electrolyte concentrations. Parfitt and Atkinson (1976) reported a a decrease of the pH for ZPC from 8.1 to 5.1 as a result of the adsorption of 100 ymole of NaH2P0^ on goethite. According to Mekaru and Uehara (1972), each millimole of P sorbed increased the CEC by approximately 0.8 meq at pH 7 on some ferruginous tropical soils. In this study, when using synthetic goethite, ZPC was at pH 5.8 for goethite (shown in Fig. 36), compared to ZPC at pH 5.2 for phosphated goethite (as a result of 3000 yg P adsorbed per g goethite). From Eq . (66) it was deduced from Gouy-Chapman theory that there should be an increase in negativity as a result of phosphate adsorption. In the

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Adsorbed (meq per 100 g Goethite ) 109 Fig. 35. Fctentiometric titration of goethite. Note point of charge occurs at pH 5.8. zero

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Sdsorbed (meq per 100 g Goethite) 110 Fig. 36. Fotentiometric titration of phosphated goethite. Not zero point of charge occurs at pFI 5.2.

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Ill present study, the goethite was freed of electrolyte by dialysis prior to ZPC determinations. This process may have had some effect on the nature of the goethite surface since loosely bound ions were removed in the process. In a soil system, all surfaces have various cations and anions present on the clay surfaces which affect both phosphate adsorptiondesorption and ion activities resulting after phosphate adsorption. From the present studies, both pH and electrolyte changes greatly affected the degree of phosphate desorption .

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CHAPTER VI SUMMARY AND CONCLUSION The low availability of phosphate ions in tropical soils is of major concern to all workers recognizing the need to increase agricultural production. In this study, conditions favoring the formation of iron oxyhydroxides such as goethite were investigated. An attempt was made to understand the mechanism through which phosphate may be bound to goethite using infrared spectroscopy. Using goethite and/or soils as supporting medium, the influence of some factors on the kinetics of phosphate adsorption and desorption were studied. The conditions for goethite preparation were investigated. Goethite was made by mixing FeCl^ and NaOH solutions and allowing the precipitates to crystallize. The identification of the presence of goethite in the suspension was made by infrared spectroscopy supported by X-ray diffraction. The infrared spectra indicated that the bands characteristic of goethite formation are at 3200 cm"^ (Fe OH stretching), 890 cm ^ and 790 cm ^ (which are both Fe -OH bending vibrations). The appearance of 2.19 8 and 2.70 ^ peaks from the X-ray diffraction also was supporting evidence that most of the components resulting from mixtures with OH/Fe = 6 were 112

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113 goethite. Both the speed and the degree of goethite crystallization increased as the OH/Fe ratio was increased. This was seen through the increase in the above band intensities. From this study, goethite was not a major component in a suspension of OH/Fe = 1.5 which was aged only for 50 hours, as reported by some workers. With OH/Fe = 1.5 in the suspension, there was need both to age for a longer period of time and to increase the temperature of the suspension to increase the formation of Fe OH bonding. The presence of phosphate at the beginning of the ageing process weakened the structure of goethite. From the data found experimentally, evidence of competition existed between the phosphate and the hydroxyl ions for bonding with iron. The increase of phosphate concentration in the suspension (P/Fe 0.032 to 3.2) induced more and more the formation of the P 0(Fe) bonds as seen by the 1000 cm~^ vibrations. The surface adsorption of phosphate on goethite was identified through the appearance of vibrational bands' in the 1200 1000 cm ^ region. The appearance of hydroxyl deformation in the same region was also identified by treatof the phosphated goethite with D 2 O. However, it was believed that the bands at 1160, 1085, and 1030 cm"^ for phosphated goethites were characteristic of phosphate binuclear bridging, (Fe0)2P00H. The existence of such type of binding is further confirmed by the increase of negative

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114 surface charge, or decrease of the point of zero charge after the phosphate adsorption. Using synthetic goethite, phosphate adsorption and desorption were affected by the initial concentration, the supporting electrolyte, the reaction time, and pH of the system. The amount of phosphate adsorbed was proportional to the initial phosphate concentration. In goethite-solution system (1 g/1000 ml), maximum phosphate adsorption was reached when the initial solute concentration was at least 10 yg P/ml. The calculated adsorption maximum, according to the Langmuir equation, increased in each case with time, the electrolyte concentration, and the valence of the cation. Increase in phosphate adsorption resulted from an increase in salt concentration. This was also accompanied by a substantial increase in phosphate desorbed by water. The increase in phosphate adsorption due to increasing cation valence was also accompanied by an increase in phosphate desorption by the water in the corresponding electrolyte over that for water only. The addition of cations might have contributed to the phosphate adsorption either through cation bridging or to a multilayer of physically adsorbed phosphate. This was supported by the fact that an increase of cation valence was not accompanied by an increase of the apparent sorption energy constant. Only with time was the phosphate adsorption reaction accompanied by an increasing

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115 sorption energy constant. Changing of the goethite suspension and phosphate systems from pH 2 to pH 12 resulted in a change in phosphate adsorption because there was competition between the hydroxyl and phosphate ions for adsorption sites. The adsorption decreased almost in a linear fashion as the pH increased. At the same initial equilibrium P concentration, the phosphate desorption was reduced when the adsorption solution pH was less than 6; but when the pH was greater than 6 the desorption phosphate increased. This showed that more phosphate was displaced from the chemically adsorbed forms to desorbable forms. This phosphate desorption increased with time up to 6 hours and remained constant from 6 to 12 hours. Both Langmuir and Freundlich equations were used to describe the relationship between the amount of phosphorus absorbed and the equilibrium concentration in goethite and soils systems. Since those equations were valid only at a particular reaction time, a kinetic model was proposed to describe the change of ions in solution as the reaction time increased. The model fractionated phosphate absorption change with time for the physically adsorbed and reversible chemically adsorbed phosphate ions compared to that for the irreversible chemically adsorbed phosphate. From the proposed •kinetic model, an expression for minimum energy for irreversible fixation was deduced. The experimental data showed

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116 that before the equilibrium state was reached, the adsorption reaction proceeded through two essentially linear-type reactions with respect to the reaction time. The phosphate adsorption isotherm for soils indicated no major differences from those of other workers. The Kenya soil had the highest adsorption compared to that of a Colorado soil and a Georgia soil. Phosphate adsorption reached the equilibrium state after 22 to 24 hours of reaction time. For any initial concentration greater than 5 yg P/ml, desorption of phosphate by 0.01 N CaCl^ solution increased with the absorption reaction time for goethite and for the Kenya and Georgia soils, but not for the Colorado soil. This desorption was in agreement with the fact that after phosphate desorption by 0.01 N CaCl 2 , further phosphate extracted by the 0.5 N NH^^F increased as the adsorption reaction time increased. Some of the phosphate adsorption at shorter time periods was extractable by 0.01 N CaCl^, but at longer adsorption reaction periods, phosphate bonding strength increased. Some of the factors determining the phenomena of adsorption-desorption have been investigated. In order to relate these studies to phosphate fertilization of tropical soils and phosphate adsorption-desorption reaction, such studies conducted in the laboratory should be complemented by pot studies involving tropical plants. Later, from these

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117 findings, field experiments would be necessary to provide demonstration plots on advances in technology for phosphorus fertilization of tropical soils.

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LITERATURE CITED Aharoni, C., and M. Ungarish. 1977. Kinetics of activated chemisorption. J. Chem. Soc . Faraday Trans I. 73:19431950. Atkins, P. W. 1978. Physical chemistry. p. 192. Freeman Company. San Francisco. Atkinson, R. J., R. L. Parfitt and R. C. Smart. 1974. Infrared study of phosphate adsorption on goethite. J. Chem. Soc. Faraday. Trans. I 70:1472-1479. Atkinson, R. J., A. M. Posner, and J. P. Quirk. 1968. Crystal nucleation in Fe(III) solutions and hydroxide gels. J. Inorg. Nucl. Chem. 30:2371-2381. Bach, B. W. , and E. G. Williams. 1971. A phosphate sorption index for soil. J. Soil Sci. 22:289-301. Baldwin, J. P., P. H. T. Nye , and P. B. Tinker. 1973. Uptake of solutes by multiple root systems from soils. Ill A model for calculating the solute uptake by a randomly dispersed root system developing in a finite volume of soil. Plant and Soil 38:621-635. Ballard, R. 1974. Extractability of reference phosphates by soil test reagents in absence and presence of soils. Soil and Crop Soc. Sci. Fla. Proc . 33:169-174. Barrow, N. J., P. G. Ozanne, and T. C. Shaw. 1965. Nutrient potential and capacity. Aust . J. Agric . Res. 16:61-76. Barrow, N. J., and T. C. Shaw. 1975. The slow reactions between soil and anions: 2. Effect of time and temperature on P concentration in solution. Soil Sci. 119:167-177 . Beckett, P. H. T. 1971. Potassium potentials. Pot. Rev. Subj . 5, 30:1-41. Beckett, P. H. T., and P. E. VJhite . 1964. Studies on the phosphate potentials of soils. Fart III. The pool of labile inorganic phosphate. Plant and Soil 21:251-281. 118

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119 Bolt, G. H. 1976. Soil Chemistry. A. Basic elements. p. 43-44. Elsevier Scientific. Publ. Co., Amsterdam. L. and G. F. Claringbull. 1965. The crystal structure of minerals. Bell Publishing Co., London. Brunauer, S. 1943. The adsorption of gases and vapors. Vol I. Physical adsorption, p. 196-217, Princeton Univ. Press. Busca, G., N. Cotena, and P. F. Rossi. 1972. Infrared spectroscopic study of micronised goethite. Mater. Chem. 3:271-284. Clark, A. 1970. The theory of adsorption and catalysis. P.258. Academic Press, Mew York and London. Colthup, N. B. , L. H. Daly, and S. E. Wiberley. 1975. Introduction to infrared and raman spectroscopy. 2nd ed. Acad. Press., New York. de Camargo, 0. A.,^J. W. Biggar, and D. R. Nielson. 1979. Transport of inorganic phosphorus in an Alfisol. Soil Sci. Soc. Amer. Proc. 43:884-890. F. A., G. B. B. Frick, and K. Sontheimer. 1978. A simplified competitive equilibrium adsorption model. Chem. Eng. Sci. 53:1667-1673. Eying, A. H. and E. M. Eying. 1963. Modern chemical kinetics. p. 10-13. Reinhold Publ. Co. New York. Farmer, V. C. and F. Palmieri. 1975. The characterization of soil minerals by infrared spectra. p. 573. In Soil components V2, Gieseking (ed.), Springer-Verlag , New York . Ghani , M. 0. and M. A. Islam. 1946. Phosphate fixation in acid soil and its mechanism. Soil Sci. 62:293-306. Hasemqn, J. F., E. H. Brown, and C. D. Whitt. 1950. Reaction of phosphate with clay and hydrous oxide of iron and aluminum. Soil Sci. 70:257-271. Hemwall, J. B. 1957. Phosphorus fixation. Adv. Agron. J: 95-112. Kingston, E. J., R. J. Atkinson, A. M. Posner, and J. P. Quirk. 1967. Specific adsorption of anions. Nature, 215:1459-1461.

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120 Hingston, F. J. , R. J. Atkinson, A. M. Posner, and J. P. Quirk. 1968. Anion adsorption by goethite and gibbsite. I the role of the proton in determining adsorption envelopes. J. Soil Sci. 23:177-191. Hingston, F. J., A. M. Posner, and J.'K. Quirk. 1974. Anion adsorption by goethite and gibbsite. II Desorption of anions from hydrous oxide surfaces. J. Soil Sci. 25:16-26. Holford, I. C. R. , and G. E. G. Mattingly. 1975. Phosphate sorption by Jurassic oolitic limestones. Geoderma, 13:257-264. Hsu, P. H. 1965. Fixation of phosphate by A1 and Fe in acidic soils. Soil Sci. 99:398-402. Hsu, P. H. 1972. Nucleation polymerization and precipitation of FeOOH. J. Soil Sci 23:409-419. Jaroniec, M, and J. Toth. 1976. Adsorption of gas mixtures on heterogeneous solid surfaces. Coll, and Poly. Sci. 254:643-649. Jossen, L. , J. M. Prausnitz, V/. Fritz, E. U. Schlunder, and A. L. Myers. 1978. Thermodynamic of multi-solute adsorption from dilute aqueous solutions. Chem. Eng. Sci. 33:1097-1106. Keng, J. C. M., and G. Uehara . 1974. Chemistry mineralogy and taxonomy of oxides and ultisols Soil Sci. Soc . Amer. Proc . 33:119-126. Kiselev, A. V., and V. I. Lygin. 1975. Infra-red spectra of surface compounds. p. 60. John Wiley and sons. New York. Kuo, S., and E. G. Lotse. 1972. Kinetics of phosphate adsorption by Ca-carbonate and Ca-kaolinite . Soil Sci. Soc. Am. Proc. 36:725-729. Laverdiere, M. R. and R. M. Weaver. 1977. Change characteristics of spodic horizons. Soil Sci Soc. Amer. Proc. 41:505-510. Lindsay, W. L. , and E. C. Moreno. 1960. Phosphate phase equilibrium in soil. Soil Sci. Soc. Amer. Proc. 24:177182.

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121 Lindsay, W. L. and P. L. G. Vlek. 1977. Phosphate minerals, p. 639. In minerals in soil environments. Soil Soc. Sci. Amer. Madison, Wis. Lingstrom, F. T. , L. Barsma, and H. Gardina. 1968. 2, 4D diffusion in saturated soils: A mathematicals theory. Soil Sci. 106:107-113. Lingstrom, F. T. , R. Haque , and V/. R. Coshow. 1970. Adsorption from solute. Ill A new model for the kinetics of adsorption-desorption processes. J. Phys . Chem. 74:495-502. Mekaru, T., and G. Uehara. 1972. Anion adsorption in ferruginous tropical soils. Soil Sci Soc. Amer. Proc . 36:296-309. Muljadi, D., A. M. Posner, and J. P. Quirk 1966 The mechanisms of P adsorption by kaolinite, gibbsite and pseudobaehmite. Part I isotherms and effect of pH on adsorption. J. Soil Sci 17:213-227. Murphy, P. H., A. M. Posner and J. P. Quirk 1975 Chemistry of iron in soils. Ferric hydrolysis products. Aust. J. Soil Res. 13:189-201. Nakamoto, K. 1978 Infrared and raman spectra of inorganic and coordination compounds. 3rd ed. John Wiley and Sons, New York. R. L. 1979. The nature of the phosphate goethite (a-FeOOH) complex formed with Ca(H„P0^)„ at different surface coverage. Soil Soc. Amer.'^ProcT 43:623-625. PL. , and R. J. Atkinson. 1976 . Phosphate adsorption on goethite (a-FeOOH). Nature. 264:740-742. Parfitt, R._L., R. J. Atkinson, and R. C. Smart 1975 The mechanism of phosphate fixation by Iron oxides. Soil Sci. Soc. Amer. Proc. 39:837-841. Parks, G. A. and P. L. de Bruyn. 1962. The zero point of charge of oxides J. Phys. Chem. 66:967-972. Peterson, G. V.’. and R. B. Corey. 1966. A modified Chang and Jackson procedure for routine fractionation of inorganic soil phosphates. Soil Sci. Soc. Amer. Proc. 30 : 563-565 .

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122 Russel, J. D., R. L. Parfitt, A. R. Fraser, and V. C. Farmer. 1974. Surface structures of gibbsite goethite and phosphated goethite. Nature 248:220-221. Ryden, J. C. and J. K. Syers . 1975. Rationalization of ionic strength and cation effects on phosphate sorption by soils. J. Soil Sci. 26:395-406. Schwertman, U., and R. M. Taylor. 1977. Non oxides. M. J. B. Dixon (ed). Minerals in soil environments, p 145180, Soil Soc. Amer. , Madison, Wis. Schofield, R. K. 1955. Can a precise meaning be given to available soil phosphorus. Soil and Fertilizers, 28:373-375. Sennett P. and J. P. Olivier. 1965. Colloid dispersion, electronic effects and the concept of zeta potential. Amer. Chem. Soc. 57:32-50. Soil investigation staff. 1972. Soil survey laboratory methods. U.S. Dept. Agr . report No. 1. Syers, J. K. , M. G. Brauman , G. VJ. Smillie, and R. B. Corey. 1973. Phosphate sorption by soils evaluated by the Langmuir adsorption equation. Soil Sci. Soc. Amer. Proc. 37:358-364. van Olphen, H. 1977. An introduction to clay colloid chemistry for clay technologists, geologists and soil scientists. 2nd ed. , John Wiley and Son, New York. Vlad, M., and E. Segal. 1979. Generalization of the Jaroniec isotherm. Surf. Sci. 79:608-616. Yuan^ T. L. , W. K. Robertson, and J. R. Weller. 1960. Forms of newly fixed phosphorus in three acid sandy soils. Soil. Sci. Soc. Amer. Proc : 24 : 44 7-45 0 .

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BIOGRAPHICAL SKETCH Yekini Bourahim was born on December 13, 1948 in Dakar (Senegal). He completed his high school education at Lycee Technique Coulibaly of Cotonou (Benin) where he obtained the degree Baccalaureat (mathemat ique ) in 1970. He immediately entered the Institut National Agronomique de Tunis and obtained after four years his Ingenieur Agricole degree. He then entered the Universite Nationale du Benin and obtained his Ingenieur Agronome (Science du Sol) degree in 1976. He served for one year as Assistant Professor at the Universite Nationale du Benin. He entered the University of Florida (Soil Science department) in 1977, and he is currently a Ph.D. degree candidate. 123

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certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a thesis for the degree of Doctor of Philosophy. John G. A. Fiskell, Chairman Professor of Soil Science I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a thesis for the degree of Doctor of Philosophy. j. j. stA set ' Science I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a thesis for the degree of Doctor of Philosophy. V. ^ Berkheiser Assistant Professor of Soil Science I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a thesis for the degree of Doctor of Philosophy. V/. K. Robertson Professor of Soil Science

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I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a thesis for the degree of Doctor of Philosophy. This dissertation was submitted to the Graduate Faculty of the College of Agriculture and to the Graduate Council, and was accepted as partial fulfillment of the requirements for the degree of Doctor of Philosophy. June 1980 Associate Professor of Chemistry Dean iculture Dean, Graduate School