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Ionic reactions in calcareous soils Forde, St. Clair McDonald 1961

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IONIC REACTIONS IN CALCAREOUS SOILS by ST. CLAIR MCDONALD FORDE B.S.A., The University of B r i t i s h Columbia, 1959 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE RETIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE IN AGRICULTURE in the Department of SOIL SCIENCE We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA April , 1961 In presenting t h i s t h e s i s i n p a r t i a l f u l f i l m e n t of the requirements f o r an advanced degree at the U n i v e r s i t y of B r i t i s h Columbia, I agree tha t the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r reference and study. I f u r t h e r agree that permission f o r extensive copying of t h i s t h e s i s f o r s c h o l a r l y purposes may be granted by the Head of my Department or by h i s representatives. It i s understood that copying or p u b l i c a t i o n of t h i s t h e s i s f o r f i n a n c i a l gain s h a l l not be allowed without my written permission. Department of The U n i v e r s i t y of B r i t i s h Columbia, Vancouver 8, Canada. Date i i ABSTRACT The s o l u b i l i t y of calcium carbonate i n calcareous s o i l s was studied with a view of t e s t i n g the v a l i d i t y of the a p p l i c a t i o n of equations based on c a l c i t e s o l u b i l i t y i n calcareous s o i l s . For t h i s purpose the s o l u b i l i t y of calcium carbonate In calcareous s o i l s was examined under a v a r i e t y of experimental conditions. Calcareous s o i l s e q u i l i b r a t e d i n water at constant p a r t i a l pressure of carbon dioxide showed a state of supersaturation with respect to c a l c i t e . The values for pH-%pCa for each s o i l were constant, but var i e d from one s o i l to another. Varying the soil-water r a t i o over a range of 1:2 to 1:10 did not a f f e c t the constancy of pH-%pCa values i n the s o i l s studied. The addition of C a C l 2 varying i n concentration from 0.0025 to 0.01 mole/1, did not a f f e c t the constancy of values f o r pH-%pCa. When d i s s o l u t i o n of part of the calcareous materials present was eff e c t e d by the addition of increments of a H-Resin or d i l u t e HCl a c i d , the values f o r pH-%pCa did not remain constant. The reactions between orthophosphate solutions and calcareous s o i l s varying i n calcium content was examined. I t was found that i n a l l cases the i n i t i a l rate of phosphorus f i x a t i o n was r a p i d with about 98% of the phosphate added being f i x e d within 2 days. The calcium carbonate exerted i t s normal e f f e c t on the s o l u t i o n equilibrium regardless of the mechanism of the r e a c t i o n between s o l i d calcium carbonate and d i l u t e solutions of KH 2P0^. The a p p l i c a t i o n of s o l u b i l i t y c r i t e r i a revealed a state of supersaturation with respect to dicalcium phosphate. A study was made to determine whether the proton transfer mechanism associated with hydrated aluminum ions was responsible for the constancy of the values for pH-%pCa in soi l s . The constancy of ion ratios was studied in cation exchange resin suspensions. It was found that the values for pH-%pCa were constant in resins saturated with calcium and hydrogen over a range of calcium saturation from 21% to 85%. At 95% calcium saturation, constancy of pH-%pCa was not observed. In a calcium-aluminum resin saturated 47% with aluminum, the values for pH-%pCa were found to be relatively constant with a maximum difference of 0.08. i v CONTENTS Page 1. INTRODUCTION _ 1 2. LITERATURE REVIEW 2 (a) The solubility of calcite in calcareous so i l s . 2 (b) Phosphate equilibria in calcareous so i l s . - . 4 (c) Ionic ratios in cation exchange resins. 5 3. DESCRIPTION OF ANALYTICAL METHODS 9 (a) Determination of pH. — 9 (b) Determination of Phosphorus. 9 (c) Determination of Aluminum. - 10 (d) Determination of Calcium. - 10 4. CALCULATION OF ION ACTIVITIES - 13 (a) Calculation of ion a c t i v i t i e s . _ 1 3 (b) Calculation of H 2P0£ and AP a c t i v i t i e s . 14 5. MATERIALS. _-_ 15 6. RESULTS AND DISCUSSION ---- 16 Part 1..SOLUBILITY OF CALCIUM CARBONATE IN CALCAREOUS SOILS.-- 16 Part 2..FIXATION OF PHOSPHORUS IN CALCAREOUS SOILS. - 28 Part 3..ION RATIOS IN EXCHANGE RESIN SUSPENSIONS - 34 7. SUMMARY. - 40 8. APPENDICES - 41 9. BIBLIOGRAPHY - 48 V LIST OF TABLES Page Table 1. The concentration of calcium in s o i l extracts as determined by flame-photometry and EDTA t i t r a t i o n . — — - — - 11 Table 2. A comparison of the relationship pH-%pCa when calcium was determined by flame-photometry and EDTA t i t r a t i o n . 12 Table 3. The location, texture and calcium carbonate content of the soils used. 15 Table 4. The solubility of calcium carbonate in some calcareous soils at P C Q^ - 0.00055 atm. 17 Table 5. Solubility of calcium carbonate in calcareous soils as affected by varying soil-water ratios at P c 0 £ - 0.00035 atm. 19 Table 6. Solubility of calcium carbonate in calcareous soils as affected by varying soil-water ratios at P C Gg z 0.00027 atm. 19 Table 7. Solubility of calcium carbonate in calcareous soils as affected by varying i n i t i a l concentrations of CaCl 2 at P C Q £ r 0.00055 atm. 21 Table 8. Solubility of calcium carbonate in calcareous soils as affected by the addition of H-Resin at P C Q 2 - 0.00039 atm. 22 Table 9. Solubility of calcium carbonate in calcareous soils as affected by the addition of dilute HCl at P C 0 £ = 0.00017 atm. 24 Table 10. Effect of KC1 on the solubility of calcium carbonate- 27 Table 11. Amount of phosphorus fixed as a function of time in calcareous soils varying in calcium carbonate content- 29 Table 12. Values for pH-%pCa in the 5%-CaC03 s o i l as a function of time. 30 Table 13. The values of pH-%pCa and pH + pI^PO^ for 'calcareous' soils of varying CaCO-j content at P c 0 2 r 0.00011 atm. 32 Table 14. The effect of increasing CaCl 2 concentration on the value of pH-%pCa for Ca-H saturated resins. 35 Table 15. The effect of increasing CaClo concentration on the value of pH-%pCa for Ga-Al saturated resins (477. saturated with Al.) 38 Table 16. Effect of Aluminum concentration on the determination of Calcium by the EDTA Method with Cal-Red the indicator 38 v i FIGURES Page Figure 1 . S o l u b i l i t y diagram for the calcium phosphates- 3 3 ACKNOWLEDGEMENTS. The author is indebted to Dr. J. S. Clark, former Assistant Professor of Soils, University of B r i t i s h Columbia, (now Pedologist, So i l Survey Branch, Canada Department of Agriculture) for his guidance and helpful suggestions during the course of this work. He expresses his appreciation also to the members of his special committee; Dr. C. A. Rowles, Professor and Chairman of the Department of So i l Science, University of B r i t i s h Columbia; Dr. J. S. Basaraba and Dr. W. D. Ki t t s , for their instruction and encouragement during his course of study at the University of B r i t i s h Columbia. The author wishes to record thanks to Mr. L. Farstad, Senior Pedologist, S o i l Survey Branch, Canada Department of Agriculture, for his help in obtaining s o i l samples and for the interest he took in the progress of this work. The author is particularly grateful for the financial assistance received from the Canada Department of Agriculture through Extra Mural Research Grant 96--"Fixation of Phosphorus in Calcareous Soils". 1 INTRODUCTION The importance of calcium carbonate in the phenomenon of phosphate fixation in soils i s now well recognized. This arises from the fact that in soils there exists a dynamic equilibrium between the phosphates in solution and the other s o i l components. Since the nature of the solid phases present w i l l influence the equilibrium conditions found at any particular time, an investigation of the solu b i l i t i e s of these solid components should provide c r i t e r i a for studying these s o i l systems. The system CaCO-j-COj-l^O therefore has an important role in calcar-eous s o i l s , when equilibria involving the basic calcium phosphates are considered. In most investigations relating to calcareous so i l s , the assump-tion i s made that the calcareous materials present react like pure calcite. On this basis, equations relating to the solubility of calcite have been applied i n the study of calcareous s o i l s . Solubility c r i t e r i a when applied to the study of systems involving the basic calcium phosphates, have made use of these same assumptions. The equilibria involving exchangeable calcium and hydrogen ions are also important in these systems. As a result of these considerations, this study was undertaken to examine: 1) the applicability of equations based on the solubility of pure calcite in calcareous so i l s , 2) the role played by calcite in phosphorus fixation, and 3 ) the factors governing the constancy of (H't)/(Ca**")^ ion ratios in so i l s . LITERATURE REVIEW 1. THE SOLUBILITY OF CALCITE IN CALCAREOUS SOILS, In the study of soils containing calcareous materials, the assump-tion i s widely made that these calcareous materials behave like pure calcite. (30, 47). Upon this basis, equations relating calcium carbonate solubility to C0 2 pressure have been developed and applied to calcareous s o i l s . (8, 11, 38, 41, 42, 44, 47). In order to derive these equations, use must be made of the solu b i l i t y product of calcite, the dissociation constants of HjCO^ and ion activity coefficients. The solubility of calcite can be represented by, Ca C0 3 = Ca 4- COj , (1.1) and aCa aC03 = K 8 p ( 1 - 2 ) Since the activity (a) of CO^ w i l l be related to the partial pressure of CO^ in the system, then account must be taken of the reactions H20 + C0 2 = H2C03, (1.3) H 2C0 3 - H + + H C O 3 , (1.4) HCO^ = H + + C0~. (1.5) The dissociation constant k^ for equation (1.4) is given by K l = SHC0 3 8H a H 2 0 C 0 2 since a = fm ; where f = activity coefficient and m = molality, then k l = mH ^ 0 3 fH F H C 0 3 ( L ' 6 > M C 0 2 F C 0 2 a H 2 0 Taking the reference state such that f = a = 1 ; then from Henry's 2 ^2® Law (12), m_ = CP__ ; where C = Henry's Law constant, and equation (1.6) V/Wj V/U2 becomes k, = a a 1 H HCOi (1.7) c p c o 2 The dissociation constant for equation (1.5) i s given by k2 = *H aC0-aHC0 3 (1.8) By substituting equations (1.7) and (1.8) into (1.2), a = ksp a % ° a H 2 (1.9) k 2 k i C P C 0 2 Rewriting equation (1.9) in terms of the notation of negative logarithms and rearranging, pH - %pCa « %(pk g p - pk 1 - pk 2 - pC - pC0 2). (1.10) Substituting the values (at 25°C) for k =4.82 x 10"9 (11) -7 k t =4.45 x 10 (13) k 2 =4.69 x 1 0 " U (14) and C = .0344 (13), equation (1.10) becomes pH - % pCa - 4.93 - % pC02 . (1.11) The equation derived above has been applied directly to calcareous s o i l s . Thus in the determination of s o i l reaction, the implications associated with the presence of CaCO-j must be observed as has been pointed out by Yaalon (48), Turner (41), and Turner and Clark (42). 4 2. PHOSPHATE EQUILIBRIA IN SOILS Interest in the basic calcium phosphates has centered about the concept that when water-soluble phosphorus compounds are added to so i l s , the majority of the phosphorus is rendered insoluble within a short time (9, 10, 28). In neutral and alkaline soils i t i s thought that relatively soluble monocalcium phosphate appears to react with calcium to give less soluble dicalcium phosphates and then highly insoluble basic calcium phosphates. Recently the use of solubility c r i t e r i a has been introduced towards solving the problem of phosphate solubility in s o i l s . Clark (6) and Lindsay et al^ (24, 25) have shown how solubility diagrams can be devel-oped and applied to soils to describe the nature of the equilibrium present. Solubility diagrams for pure calcium phosphates show a linear relationship between the "monocalcium phosphate potential" (pHjPO^ + % pCa), and the "lime potential" (pH-%pCa) ( 6 ) . A better way of describing the system would be in terms of pH-%pCa and (pH + p^PO^). A solubility diagram for the pure calcium phosphate compounds is shown in Appendix 1. The lines are based on values given for the solubility products for the calcium phosphates by Lindsay (24). The equations were developed as follows: For the general case, the formation of the basic calcium phosphates may be represented:-x Ca (OH)2 + y H^ PO^  z Calciumphosphate + w H20. (2.1) The equation in terms of the chemical potential of the reacting species and the products formed, is given by x « C a ( 0 H ) 2 + % 3 P 0 4 = Z ^ C a P ( 2 ' 2 ) and x <P°Ca(0H)2 + * T 1« aCa(0H)2) + 70*°*^ + R T l n aH 3P0 4> = 2 *CaP' <2' 5 Rearranging (2.3) x R T l n aCa(OH) 2 + y RT In z k. (2.4) where k = z#CaP - ^°ca(OH) 2 " ^ P C * ' The activity of a completely dissociated electrolyte is the product of the activity of the constituent ions, thus, aCa(OH) 2 = a C a *oi/» and aCa(OH) 2 = p C a + 2 p 0 H -After dividing by 2.303 RT and expressing the negative logarithm of the activity in accordance, with the convention of pH, equation (2.4) becomes, x (pCa + 2 pOH) + y (pH + pHjjPO^) = -k' . Simplifying the above, -2x (pH - %pCa)+ y (pH+ pH 2P0 4) = k". (2.5) Larsen (21) found that when different rates of a soluble phosphate source were added to s o i l samples and equilibrated for some time, a plot of the "monocalcium phosphate potential" (pH2P0^ + %pCa) verus the "lime potential" (pH -%pCa) gave points which ranged from a condition of undersaturation with respect to hydroxyapatite to supersaturation with respect to dicalcium phosphate. 3. IONIC EQUILIBRIA IN SOILS The relationship between the concentrations of two different ions present in any solution in equilibrium with the same s o i l sample was f i r s t investigated by Terasvuori (40). He argued that on the basis of the Donnan equations, the ratio (H**)/(Ca)^ should have a constant value in any solution 6 in equilibrium with a s o i l containing calcium as the main exchangeable cation. This he showed to be approximately true by experimental methods. A more general examination of the relation between the concentra-tions of a l l the ions present in the equilibrium solutions was made by Schofield (34, 36). He interpreted the relationship between the concentra-tions of a l l ions present in the equilibrium solutions obtained from percolation through various s o i l samples, on the basis of the "ratio law" derived from the Gouy theory of the el e c t r i c a l double layer. From the theory of the electric double layer i t i s assumed that ions "evaporate" from the surface of constant surface charge density, giving r i s e to a diffuse distribution as a result of their kinetic energy. These ions satisfy charges on the surface. The Coulomb interaction between the charges present in the system can be described by the Poisson's equation, * V - -—111 ( 3 . 1 ) D where + = the potential and describes the change in potential from a certain value f oat the interface to zero in the bulk solution. A = the Laplace operator. f ° net space charge density. J> = dielectric constant. The distribution of ions in solution is governed by the Boltzmann relation (19) expressing that at places of positive potential the negative ions are concen-trated and the positive ones repelled, whereas for places of negative potential the reverse is the case. Thus the concentration of ions in the bulk solution in equilibrium with the clay particles may be related to the concentration of the respective ions on the clay surface by the Boltzmann's equation which, when applied to both cations and anions becomes: z* z+ (M ) = (M ) exp (-Ze*7kT), s o ( A Z " ) g - ( A Z " ) o exp (-ZetykT), (3.2) (3.3) in which M and A refer to some cation and anion whose valences are and Z~, respectively; e is the electronic charge; f is the electric potential on the surface of the clay particle; k is the Boltzmann constant; T is the absolute temperature; the subscripts s and o refer to the surface of the clay particle and the bulk solution, respectively; and parentheses indicate ion a c t i v i t i e s . Combining equations (3.2) and (3.3) and remembering that Z-in equation (3.3) has a negative sign, or (M Z +) ( l/Z z + i/z l/z" (A" ), (A 2"), l/Z" constant (3.4) z + l/Z + z . l/Z (M ) (A ) (3.5) When applied to the more general case to two cations and M^, the relation-ship given by equation (3.4) leads to + zt u** (M^l) 7+ 1/Z+ Zt. 1/ 2! (M^l) z + i / z , + (M2 2) These equations would predict that for an exchanger having the surface concentrations, the reduced ratio remains constant through the double-layer. Applying this to a situation involving H** and Ca"1-*, we have + oo. constant (3.6) / ( C a ) 8 Thus i f the external concentration i s not changed markedly the r e s u l t i s (H +) = (H +) = constant (3.17) 7(Ca++) / ( C a ^ ) before a d d i t i o n a f t e r a d d i t i o n of e l e c t r o l y t e of e l e c t r o l y t e Calcium i s the predominant cation i n humid s o i l s making up about 50% of the exchangeable cations. I f CaCl2 i s added i n varying concentrations, since the surface concentration i s constant the r a t i o of surface concentra-t i o n to external concentration should remain constant. As far as hydrogen i s concerned most data i n d i c a t e that there i s not much exchangeable hydrogen i n inorganic c l a y s . (1, 26, 33, 35) However, the d i s s o c i a t i o n of A 1 + 3 ( H 2 0 ) 6 (3), A 1 + 3 ( H 2 0 ) 6 = A1(0H)(H 20)$++H +, (3.8) would provide a proton t r a n s f e r mechanism that may be responsible f or keeping the surface concentration of (H ) constant i n s o i l s . 9 DESCRIPTION OF ANALYTICAL METHODS a) DETERMINATION of pH The determination of pH proved to be the most c r i t i c a l measurement in these studies, and especial care was exercised to ensure that reliable pH measurements were obtained. The instrument used was a Cambridge pH meter of the bench type equipped with a wide range glass electrode and a Leeds Northrup calomel reference electrode. The instrument was standardized using buffer solutions of pH 4, 7 and 9 (51). The methods for making up these buffer solutions are described in Appendix I I . The buffer solution of pH 4, gives a pH of 4.00 at 15°C and has a temperature coefficient of p Ht°c = 4 ' 0 G + °' 5 (t-15/100) 2 for o o t between 0 C and 60 C. The buffer solution of pH 9, gives a pH of 9.27 at 15°C and has a temperature coefficient of p Ht°c = 9 , 2 7 " ° ' 0 0 8 5 <t " 1 5> f o r t between 0°C and 60°C. In the solubility studies, a i r at the same partial pressure of C0 2 as that used to equilibrate the suspensions during the experiments, was bubbled through the suspensions when pH measurements were made. The suspensions were placed in polyethylene containers while pH measurements were taken. b) DETERMINATION OF PHOSPHORUS Phosphorus was determined using the molybdenum blue method described by Jackson (16). After allowing the colour to develop for a period of four 10 minutes, the per cent transmission was read on a Bauch and Lomb colorimeter at 660 mu. The details of the method are given in Appendix 111. c) DETERMINATION OF ALUMINUM Aluminum was determined by the Aluminon method (23). Details of the method are given in Appendix IV. d) DETERMINATION OF CALCIUM CONCENTRATION (1) The Flame- photometric method compared with EDTA titrations. The solutions on which calcium and magnesium determinations were made, were extracts obtained after f i l t r a t i o n of s o i l suspensions that had been e q u i l i -brated for several days. Of the many methods available for the determination of the concentrations of calcium and magnesium, that by flame- photometry and others by EDTA titrations were considered. Since the possibility existed that some of the calcium present in the s o i l extracts may have been in a complexed state (30), i t was decided to compare the analysis for calcium as determined by flame-photometry with that using EDTA titrations with Cal-Red as indicator (32). This latter method, details of which are given in Appendix V makes use of the indicator-2 hydroxyl-l-(2 hydroxy -4-sulfo-l-napthylazo)-3-napthoic acid-which permits t i t r a t i o n of calcium in the presence of magnesium with EDTA at pH 12. Such analyses were made on the extracts obtained from the s o i l samples after f i l t r a t i o n . These extracts w i l l be referred to as "untreated". To aliquots of the f i l t r a t e , 10 ml. of 30% H202 and 5 ml. of concentrated HCl were added and the solution allowed to evaporate to dryness on a steam bath and then made up to volume. Calcium concentrations were determined by flame-photometry as well as by EDTA titrations using the Cal-Red indicator. These solutions to which H 20 2 and HCl were added w i l l be referred to as "treated". 11 Table 1. The concentration of calcium in s o i l extracts as determined by flame-photometry and EDTA titration: Calcium Concentration in moles/1 x 10 Flame Determination EDTA Determination Percentage Increase Sample Treated over Untreated untreated treated untreated treated flame Cal-Red Lumby 1 1.95 2.25 1.491 1.571 15.38 5.36 2 2.05 2.25 1.573 - 24.39 -Penticton 1 0.675 0.85 0.554 0.587 25.92 6.53 2 0.650 0.85 0.554 0.587 30.77 6.53 Hilton 1 1.31 1.515 0.9576 0.987 15.65 -1.12 2 1.40 1.660 1.0093 1.045 14.28 5.57 Davis 1 0.72 0.81 0.722 0.774 12.50 7.2 2 0.72 0.81 0.722 0.774 12.50 7.2 It can be seen from Table 1, that treatment of the extracts with H 20 2 and HCl resulted in an increase in the calcium concentration when determined by both flame-photometry and the EDTA method. This would indicate that treatment of the extracts leads to a breakdown of unstable calcium complexes. The v a r i a b i l i t y of the data as shown by a comparison between duplicates for each s o i l sample would also suggest that the amount of calcium complexed varies considerably from one extract to another for the same s o i l sample. The concentrations of calcium as determined by flame-photometry were in a l l cases greater than those obtained by titra t i o n , whether the extracts were treated or untreated. This would indicate that the increase in concentration was due largely to the breakdown during flame-photometry of some of the less stable calcium-complexes. Since different values for calcium 12 concentration were obtained by using the two methods described above, i t was necessary to see whether this difference in calcium concentration affects the calculation of the relationship pH-% pCa. Values for pH-% pCa were calculated from the data in Table 1, and are shown in Table 2. It can be seen that when comparison is made of these values for pH-% pCa between the "untreated" and "treated" of any one method, the difference is not significantly different. However, the differences tend to be greater when comparisons are made between the methods themselves. Since the flame-photometric determination gave greater percentages increase in calcium concentration for the "treated" over the "untreated" extracts than did the EDTA method, i t would seem that more of the complexed calcium is determined by flame-photometry than by the EDTA method. When this was taken into consideration i t was decided to use the EDTA tit r a t i o n with Cal-Red the indicator on the untreated extracts for the determination of calcium concentrations. Table 2. A comparison of the relationship pH-%pCa when calcium was determined by flame-photometry and EDTA ti t r a t i o n . pH - % pCa values Sample Flame Determination EDTA Titration Untreated Treated Untreated Treated Lumby 1 6.71 6.74 6.66 6.67 2 6.71 6.75 6.65 -Penticton 1 6.62 6.57 6.57 6.59 2 6.67 6.63 6.63 6.65 Hilton 1 6.63 6.66 6.57 6.56 2 6.70 6.63 6.63 6.54 Davis 1 6.59 6.56 6.59 6.60 2 6.59 6.56 6.59 6.60 13 CALCULATION OF ION ACTIVITIES In order to express the concentration of a given element in terms of the activity of a particular ion species, i t was necessary to make use of ion activity coefficients and dissociation constants. The ion activity coefficients were calculated from the second approximation of the Debye-Huckel equation (19). - l o g f t z Azf Vi 1 i - B a ^ / l where: = activity coefficient of ion i = valency of ion i I = Ionic strength a^ - effective diameter of the ion in solution A and B = constants Values used for these constants were (19), T°C A in A° BxlO" 8 15 0.5000 0.3260 20 0.5042 0.3273 23 0.5060 0.3280 25 0.5085 0.3281 a i has to be specified for each particular ion and the following values were used (18), ion a^(A°) CaH* 6 Mg"*4" 8 H2P04 4 HP0| 4 POf 4 The ionic strength of the solutions was calculated from the relationship; 2 1 2, k c.i Zi» where I = ionic strength, c i= molar concentration of any ion i , Z^= valency of ion i . In calculating the acti v i t i e s of the ionic species in these studies, i t was assumed that hydrogen, calcium and magnesium (where present) were the pre-dominant cation species and that monovalent anions were present in equiva-lent concentrations. CALCULATION OF H2P0^ AND A l + 3 ACTIVITIES 4 .3 The methods used for calculating the activities of HgPO^ and Al are outlined in Appendix VI. 15 MATERIALS 1. SOLUBILITY OF CALCIUM CARBONATE IN CALCAREOUS SOILS For the studies on the solubility of calcium carbonate in soils, four calcareous soils from different parts of the Province of B r i t i s h Columbia were used. The location, texture and carbonate content (37) of the four soils are summarized in Table 3. Table 3. The location, texture and calcium carbonate content of the soils used. S o i l Series Location Texture Percent Calcium Carbonate Lumby Okanagan s i l t loam 1.91 Hilton Okanagan s i l t y clay loam 1.96 Penticton Okanagan s i l t loam 2.37 Davis Peace River sandy loam 8.48 For the study involving the fixation of phosphorus as a function of calcium carbonate content, the s o i l sample used was Alcan from the Peace River d i s t r i c t of B r i t i s h Columbia. This s o i l has a reaction of pH 4, and a cation exchange capacity of 20 me/100 grams. 16 RESULTS AND DISCUSSION Part 1. SOLUBILITY OF CALCIUM CARBONATE IN CALCAREOUS SOILS In the investigation of the solubility of calcium carbonate in calcareous so i l s , 20 gram samples of air-dried soils were placed in 250 ml. erlenmeyer flasks and 100 ml. of d i s t i l l e d water was added. The suspensions were shaken vigorously and atmospheric air was bubbled through the suspensions. After a period of 7 or 9 days, pH values were determined and the suspensions f i l t e r e d . The concentrations of calcium plus magnesium were determined on the f i l t r a t e s by use of the EDTA method with Eriochrome Black T as the indicator (52), and the concentration of calcium determined by EDTA t i t r a t i o n with Cal-Red indicator. Values for pH-% p Ca were calculated, and these are shown in Table 4. It can be seen from Table 4 that pH values varied from one s o i l sample to another, as well as did the calcium concentrations. The relation-ship pH-% pCa therefore differed from one s o i l sample to another. 17 Table 4. The solu b i l i t y of calcium carbonate in some calcareous soils at P C Q 2 : 0.00055 atm. Sample pH at 25°C (Ca) m/1 x 10 3 pCa pH-% pCa Penticton 1 8.19 0.850 3.1798 6.60 2 8.23 0.615 3.3097 6.58 3 8.24 0.730 3.2383 6.62 4 8.27 0.675 3,2700 6.64 5 8.26 0.915 3.1466 6.59 Hilton 1 8.22 1.365 2.9805 6.73 2 8.20 1.385 2.8747 6.77 3 8.19 1.385 2.8747 6.76 4 8.20 1.420 2.9654 6.72 5 8.19 1.300 2.9991 6.70 Lumby 1 8.14 2.100 2.8429 6.72 2 8.14 2.100 2.8424 6.72 3 8.14 2.100 2.8426 6.72 4 8.17 1.900 2.8785 6.73 5 8.14 2.100 2.8410 6.72 If the assumption i s made that the calcareous materials present in soils behave like pure calcite, then i t would be expected that the equilibrium value pH-% p Ca would be constant, and that this value should correspond to that obtained for pure calcite at the same partial pressure of C0 2. However, when calcareous s o i l were equilibrated in water for a period extending from 7 to 9 days i t was found (Table 4) that a state of supersaturation with respect to calcite existed, to the extent that there was a 2 to 3 fold excess of calcium in solution relative to that given by calcite at the same partial pressure of C0 2. As shown in Table 4, the values for pH-% pCa though constant for individual s o i l s , differed from s o i l to s o i l . The results shown here are similar to those reported by Olsen (30). Olsen explained the excess in calcium 18 concentration in terms of the presence of a more soluble calcareous material. The poss i b i l i t y of the existence of calcareous materials with so l u b i l i t i e s quite different from that of pure calcite was not considered. To determine whether or not a possible state of supersaturation could exist, or whether a more soluble solid phase of calcite could determine the f i n a l equilibrium concentrations the effect of varying the soil-water ratio on the solubility of calcium carbonate in soils was studied. Samples of the calcareous soils were prepared with varying soil-water ratios, and after equilibration for 7 days, pH, calcium and magnesium concentrations were determined as outlined above. It i s seen from Table 5, that in a l l cases, as the soil-water ratio decreased, the calcium concentration increased. The pH increased about 0.1 a/ unit with„dilution factor of 2.5. The pH-%pCa values while differing among soil s , were again constant for each s o i l . The constancy of the relationship pH-%pCa suggested that the excess calcium concentration was not due to a more soluble phase of calcite, but that the calcareous materials present behaved differently from calcite and had unique s o l u b i l i t i e s . The differences of -0.05 in the values of pH-%pCa were considered to be within the limits of error in the experimental methods used, particularly that of pH. 19 Table 5. Solubility of calcium carbonate i n calcareous soils as affected by varying soil-water ratios at P r o => 0.00035 atm. Sample No. Soil-.Water pH at 25°C (Ca)m/1 x 10 3 pCa pH-%pCa Lumby IA 1:10 8.32 0.947 3.1525 6.74 IB 1:10 8.32 0.947 3.1525 6.74 2A 1:5 8.24 1.561 2.9723 6.75 2B 1:5 8.24 1.561 2.9723 6.75 3A 1:2 8.10 3.333 2.7055 6.75 3B 1:2 8.14 3.188 2.7222 6.78 Penticton IA 1:10 8.40 0.471 3.4145 6.69 IB 1:10 8.40 0.471 3.4145 6.69 2A 1:5 8.39 0.523 3.3741 6.70 2B 1:5 8.39 0.479 3.4089 6.80 3A 1:2 8.26 0.826 3.2018 6.66 3B 1:2 8.26 0.921 3.1608 6.70 Table 6. Solubility of calcium carbonate i n calcareous soils as affected by varying soil-water ratios at Pc(>2 ' = 0.00027 atm. Sample No. Soil-.Water pH at 25°C (Ca)m/1 x 10 3 pCa pH=%pCa Hilton IA 1:10 8.31 0.6780 3.2681 6.68 IB 1:10 8.31 0.6289 3.2950 6.67 2A 1:5 8.24 0.9110 3.1513 6.66 2B 1:5 8.23 0.9368 3.1420 6.66 3A 1:2 8.14 1.8556 2.8848 6.70 3B 1:2 8.14 1.9694 2.8630 6.71 Davis IA 1:10 8.31 0.5564 3.3438 6.64 IB 1:10 8.32 0.5331 3.3590 6.64 2A 1:5 8.31 0.5926 3.3382 6.65 2B 1:5 8.33 0.5823 3.3250 6.67 3A 1:2 8.23 0.7013 3.2542 6.60 3B 1:2 8.28 0.7019 3.2506 6.66 20 A further examination was made of the system CaG03"H20 - C©2 by investigating the solubility of calcium carbonate in calcareous soils in the presence of varying concentrations of CaC^. Air-dried samples of the calcareous soils were equilibrated in solutions of CaCl2 varying in concentration from 2.5xl0" J M/L to 1.0x10 M/L, and analyses carried out on the f i l t r a t e s as before. The results of these experiments are shown in Table 7. It can be seen that an increase in the i n i t i a l concentration of CaCl2 resulted i n a decrease i n the pH to the extent of O.lunit for each two-fold increase i n CaClj concentration. The fin a l calcium concentration increased with an increase in the i n i t i a l CaCl2 concentration. In other words, the fin a l calcium concentration increased as the pH decreased. However, the value for pH-%pCa was maintained constant for a given s o i l even though the i n i t i a l concentration of CaCl2 was allowed to vary from 2.5 x 10~3 M/L to 1.0 x IO" 2 M/L. Again, this data would suggest that the calcareous materials behave differently from pure calcite. A comparison of the values for pH-%pGa for the two soils used'here, with those for pure calcite, would show a state of supersaturation with respect to calcium carbonate in the soil, solutions. 21 Table 7. Solubility of calcium carbonate in calcareous soils as affected by varying i n i t i a l concentrations of CaC^at P^C^ ^ 0.00055 atm. Sample No. I n i t i a l (Ca) m/1 x 10 3 pH @ 24°C (Ca)m/lxl0 3 pCa pH-%pCa Lumby 1A 2.5 8.04 2.9334 2.7457 6.67 B 2.5 8.04 2.9234 2.7457 6.67 2A 5.0 7.99 4.6588 2.5886 6.68 B 5.0 7.97 4.6588 2.5886 6.68 3A 10.0 7.88 8.1814 2.3797 6.69 B 10.0 7.89 8.1814 2.3797 6.69 Hilton 1A 2.5 8.07 2.5153 2.7808 6.68 B 2.5 8.07 2.5153 2.7808 6.68 2A 5.0 7.96 4.5349 2.5666 6.68 B 5.0 7.96 4.5349 2.5666 6.68 3A 10.0 7.86 8.8889 2.3324 6.69 B 10.0 7.86 8.8889 2.3324 6.69 Calcite 1A 2.5 7.97 2.9131 2.6963 6.62 B 2.5 7.92 3.3469 2.6453 6.60 2A 5.0 7.85 5.3819 2.4756 6.61 B 5.0 7.85 5.3819 2.4756 6.61 3A 10.0 7.75 10.0201 2.2617 6.62 B 10.0 7.75 10.0201 2.2617 6.62 The data in Tables 4-7 show that the calcareous materials present in soils behave differently from pure calcite. On the other hand, the constancy of the relationship pH-%pCa was not affected by varying the soil-water ratio or the i n i t i a l CaCl2 concentration indicating that the calcareous materials in each of these soils has a unique solubility. Olsen (30) has suggested that the surface of the calcium carbonate in soils may be coated with a more soluble surface phase. The solution equili-brium would then be controlled by this soluble surface phase accounting for the observed state of supersaturation with respect to calcite. An attempt was made to remove this surface phase by adding increments of a H-saturated cation exchange resin to the soils and observing the pH-%pCa values as increasing 22 amounts of the s o i l calcium carbonate was dissolved. In 1:5 soil:water ratio, Amberlite IRC-50(H+) resin was added i n amounts calculated to dissolve from 0 to 75% of the calcium carbonate i n the s o i l . After equilibration for two weeks, while atmospheric air was bubbled through the suspensions pH measurements were taken and calcium and magnesium concentrations determined on the f i l t r a t e s . The results obtained for the Lumby so i l are shown i n Table 8, together with values for pH-%pCa. Table 8. Solubility of calcium carbonate i n calcareous soils as affected by the addition of H-Resin at P C 0 2 = 0.00039 atm. Sample meq. of H-Resin added pH @ 21°C. . pCa pMg. pH-%pCa 1A 0 8.11 2.9517 2.8961 6.64 B 0 8.11 2.9525 2.9087 6.64 2A 2 7.85 2.8749 2.8936 6.41 B 2 7.87 2.8878 2.8754 6.43 3A 4 7.72 2.8469 2.8917 6.30 B 4 7.76 2.8581 2.8958 6.33 4A 6 7.65 2.7993 2.8927 6.25 B 6 7.75 2.8426 2.9068 6.23 5A 10 7.50 2.7769 2.8941 6.11 6A 15 7.26 2.7415 2.8997 5.89 B 15 7.22 2.7305 2.9219- 5.86 The results show that there was a decrease in pH with increasing additions of H-Resin. The calcium concentration increased with increasing amounts of H-Resin added, but the concentration of magnesium was not affected to any extent by increasing the amount of H-Resin added. However, the relationship pH-%pCa 23 did not remain constant but decreased with increasing additions of milliequiva-lents of H-Resin. This would mean that the calcareous material present was becoming more soluble with respect to calcite. The reactions involved with the addition of a H-Resin to a s o i l containing calcium carbonate should be: CaC03 = Ca"14" + CO3 , H-Resin + Oa** =. Ca-Resin + H+. An increase in the amount of H-Resin would cause the dissolution of CaCG-3, and the release of if from the resin. When account is taken of the reactions: H20 + C0 2 = H 2C0 3 ; H 2C0 3 ' - H + + HCO3 ; HCO§ - H* + C0| ; i t can be seen that the partial pressure of C0 2 w i l l determine the pH. In this experiment the decrease i n pH cannot be explained easily in terms of the reactions shown above. It would seem that addition of H-Resins to a calcareous s o i l does not result i n as simple an equilibrium situation as the reactions above would predict. In another attempt to elucidiate the problem, samples of calcareous soils were dispersed in d i s t i l l e d water for 9 days, after which time dilute HCl was added so as to effect dissolution of from 0 to 50% of the calcium carbonate present. These suspensions were a l l brought to a fin a l soil-water ratio of 1:5 and equilibrated for another 5-day period and then analyses were carried out as described above. Results are shown i n Table 9. As before, the pH decreased with increasing additions of dilute HCl, but the calcium concentra-tion increased. For the Hilton s o i l , pH-%pCa increased and tended to approach the value pH-%pCa » 6.82 for calcite at this partial pressure of C0 9. 24 Table 9. Solubility of calcium carbonate i n calcareous soils as affected by the addition of dilute HCl at ?CQ2 = 0.00017 atm. m l > o f .Q42N Sample No. HCl added. pH @ 25°C. pCa pH-%pGa Hilton IA 0 8.11 3.0599 6.58 B 8.16 3.0860 6.62 2A 10 8.00 2.7287 6.64 B 8.00 2.7250 6.64 3A 25 8.01 2.5080 6.76 B 8.07 2.5084 6.82 4A 50 7.91 2.3083 6.76 B 8.03 2.2988 6.88 Penticton IA 0 8.14 3.1434 6.57 B 8.14 3.1438 6.57 2A 10 8.00 2.8281 6.59 B 8.00 2.8281 6.59 3A 25 7.79 2.5939 6.49 B 7.79 2.5918 6.49 4A 50 7.66 2.3463 6.49 B 7.66 2.3463 6.49 For the Penticton s o i l an anamalous behaviour was observed. The value for pH-%pCa decreased with increasing additions of HCl and indicated a state of supersaturation with respect to calcite to exist at this partial pressure of C02* ^ n e °f t^-e chief d i f f i c u l t i e s in evaluating these results is inherent i n the experimental techniques employed here. It was assumed that the HCl would react primarily with the surface of the calcareous materials present i n the soils, and leave the crystal unaffected. However,this assumption i s not well founded since the reaction of the HCl with the calcareous material need not follow this predicted manner. It could be that the reaction would take 25 place on the surface of some of the larger calcite crystals, or i n other cases disrupt the crystals so as to complicate the final equilibrium situation found. Then too, many other solution equilibria may be affected by the addition of acid. The observed state of supersaturation i n the calcareous soils may be attributed to .differences i n the composition and solubility of the carbonate material i n the s o i l . The presence of such material with a fixed and unique solubility would account for the constant pH-%pCa values obtained with changing soil-water ratios. The presence of a small amount of a more soluble calcium carbonate, under these circumstances, does not provide a satisfactory explanation for the supersaturation of the soils with respect to calcite. Of the forms of calcium carbonate described by Brooks et a l (4), calcite, aragonite, vaterite, monohydrate and hexahydrate, could a l l be present in soils (30), but the latter three have not been reported. Vaterite has been clearly recognized as a distinct form of calcium carbonate which can be obtained by precipitation in the laboratory. Cole (8) reported having produced vaterite crystal species when CaC03 was precipitated i n the presence of small amounts of magnesium. It should be observed that many of these carbonates are metabstable in nature. Aragonite appears to be less stable than calcite at ordinary pressures and temperatures. The difference i n s t a b i l i t y between calcite and aragonite at ordinary temperatures i s quite small, however, and neither variety of calcium carbonate shows any tendency to change into the other i n the dry state at room temperature. However, McKenzie (27) quotes Kendall's data of 0.0143g/L for calcite and 0.0153 g/L for aragonite in water at 25°C under similar but not accurately known conditions of C02 partial pressure to show the differences in solubility in water. From these figures a change from aragonite to calcite i n contact with water in a s o i l can be anticipated. 26 Vaterite i s said to be less, stable (4) than calcite or aragonite. It slowly changes to the calcite modification when in the moist condition. Since i t has a higher solubility than calcite, a condition of supersaturation with respect to calcite w i l l be observed should this species be present in a s o i l . Thus i t has significance in the study of calcareous s o i l s . It i s doubtful i f CaCC>3.6H20 can persist as such in soils, except perhaps in a metastable condition. According to Brooks et a l (4) the solubility of CaC03.6H20 is two to three times that of calcite, while the other forms are intermediate i n solubility compared to calcite. Upon this basis then i t can be concluded that the observed state of supersaturation i n calcareous soils may be due to some of these carbonate species i f present. Another factor that has not been fu l l y investigated i s that of the role played by magnesium in calcite solubility. Brooks et al (4) have shown i n a qualitative study that the presence of additives such as magnesium or 'Calgon' (a glassy form of a sodium phosphate) leads to a modification i n the crystal habit of calcite. Distortion of the calcite lattice is consistent with the formation of metastable phases which are more soluble than calcite. Further study i s also needed in regard to the solubility of calcium carbonates that may be associated with varying amounts of MgCC-3. In this connection much work has been done with regard to the system CaC03-MgC03~ CO2-H2O (39, 49, 50) i n pure systems, but applications have not been made to so i l s . Since the solubility of some crystalline compounds in soils may be influenced by various salts (4, 31) the effect of the salt KC1 on CaC03 solubility was examined experimentally. Five gram samples of calcium carbonate were placed in 250 ml. erlenmeyer flasks and dispersed i n 100 ml. of KC1 varying i n concentration from 0.0005 to 0.1 moles/L. These suspensions were 27 equilibrated at constant temperature in a water bath held at 25°C. After two weeks the calcium concentrations were determined on the f i l t r a t e s . It was very d i f f i c u l t to measure pH in these unbuffered solutions, and these measurements were omitted. The results of this experiment are shown i n Table 10. Table 10. Effect of KC1 on the solubility of calcium carbonate. KC1 Ca . PfCa pCa m/L m/LxlO 0.0005 6.17 0.400 3.30 0.0010 6.64 0.322 3.28 0.0050 7.67 0.184 3.29 0.0100 8.16 0.172 3.30 0.0500 10.45 0.101 3.30 0.1000 11.98 0.089 3.32 As can be expected, the calcium concentration increased with an increase in the concentration of the KC1, but the activity coefficients decreased with an increase i n ionic strength, so that the value for pCa remained constant. While these results would predict that the solubility product of CaC03 w o u l d not be affected by dilute salt solutions in soils, Brooks et a l showed that the presence of magnesium or phosphorus may affect calcite solubility. In view of a l l that has been said above, i t i s agreed (30) that calcareous materials present i n soils do not react like calcite since they do not a l l have the same solubility as pure calcite. As a result, equations based on calcite solubility cannot be assumed safely to apply to any calcareous s o i l . 28 Part 11. FIXATION OF PHOSPHORUS IN CALCAREOUS SOILS Use is often made of thermodynamic c r i t e r i a for interpreting phosphate reactions i n soi l s . In the study of the fixation of phosphorus i n calcareous soi l s , use is made of the properties of the pure crystalline calcium phosphates. The identification by Lehr et al(22) of CaHPO^^H^O (dicalcium phosphate dihydrate), CaHPO^ (dicalcium phosphate anhydrous), Ca^H^O^^. 3^0 (octocalcium phosphate) and apatite in soi l s , after the application of superphosphate, has removed most of the limitations involved from the theoretical treatment of phosphate reactions i n calcareous s o i l s . Since a study of the solubility of calcareous soils had revealed some complications as far as calcite solubility was concerned, i t was decided to eliminate this d i f f i c u l t y , by working with a s o i l in which the calcareous material was well defined. Toward this end, samples of an acid s o i l , (Alcan) were brought to various levels of calcium carbonate content by the addition of weighed amounts of CaCO^ to the s o i l . This s o i l had a reaction of pH 4.7 and a cation exchange capacity of 20 me/100 gm. (1) Rate of phosphorus fixation. Preliminary experiments were conducted to establish the rate at which phosphorus was fixed in the a r t i f i c i a l l y prepared calcareous s o i l s . To 20 gm. samples of s o i l , 100 ml of solution containing 47.5 micrograms of phosphorus were added as RH2PO4. Air was bubbled through the suspensions which were allowed to shake continuously. Samples were removed at intervals, and pH recorded. The suspensions were then f i l t e r e d and the f i l t r a t e s analysed for calcium and phosphorus. 29 From Table 11, i t is readily seen that fixation of phosphorus was very rapid over the i n i t i a l two day period of the experiment. Table 11. Amount of phosphorus fixed as a function of time in calcareous soils varying in calcium carbonate content. Time i n days Phosphorus fixed in micrograms 57. GaC03 107. CaC03 15% CaC03 2 46.686 46.896 46.686 7 46.973 47.087 47.087 15 47.047 — — 27 47.100 47.034 47.234 dn a l l cases, over 98% of the phosphorus added i n i t i a l l y was removed during this period. The phosphorus continued to be fixed even after 27 days, at which time, the experiment was discontinued. The result was that while the absolute amount of phosphorus fixed increased with respect to time, the rate of fixation was most rapid in the i n i t i a l period of the experiment and then decreased t i l l equilibrium was reached. Slimilar results have been reported; (2, 15, 17) in the fixation of phosphorus in calcareous soils and i n clays. For the 5% CaC03 series, values for pH-%pCa were calculated, and these are shown in Table 12. 30 Table 12. Values for pH-%pCa in the 57.-CaC03 s o i l as a function of time. Days Sample pH @ 24°C Ca o pCa pH-%pCa m/L x 10 J 2 IA 7.95 1.29 3.001 6.45 B 7.92 1.47 2.949 6.45 7 2A 8.00 1.85 2.865 6.57 B 8.03 1.57 2.926 6.57 15 3A 8.15 1.55 2.932 6.68 B 8.15 1.42 2.965 6.67 27 4A 8.15 1.18 3.036 6.63 The data in Table 12 would suggest that after 15 days equilibrium with respect to the solubility of CaC03 was being attained. It can be assumed that the values for pH-%pCa were approaching 6.66 for pure calcite i n equili-brium with carbon dioxide at 0.0003 atmosphere partial pressure. It should be noted that since magnesium was not determined in these experiments, a slight error i s introduced in the calculation of the ionic strength and hence the calculated activities of calcium may be in slight error. However,this experiment was of an investigatory nature and served the purpose of establishing the rate at which equilibrium was attained. ( i i ) The reactions between calcareous soils and orthophosphate solutions. In a c r i t i c a l experiment designed to investigate phosphorus fixation as a function of the calcium carbonate content, s o i l samples were brought to calcium carbonate contents of 2, 5, 10, and 15 per cent. 20gm samples of so i l were dispersed i n 100 ml. of KH2p04 solutions while a i r at PCO2 w a s bubbled through the suspensions, which were shaken at certain times during the day. The rate at which phosphorus was added was equivalent to 20001b/acre of 31 ?2®5' After equilibration for 36 days, pH measurements were made and calcium, magnesium and phosphorus determined on the solutions obtained by f i l t r a t i o n . In Table 13 are shown the calculated values of pH-%pCa and pH + pl^PO^. As shown by the values for pH-%pCa in the soils in which the calcium carbonate content was 2 and 5 percent, there was a state of supersaturation with respect to calcite. However in soils containing 10 and 15 percent calcium carbonate, the calcium carbonate exerted i t s normal effect on the solution equilibrium regardless of the mechanism of the reaction between solid calcium carbonate and dilute solutions of phosphate. In these two latter cases the value of pH-%pCa approached that for pure calcite at Pc02 = °'000Hatm« This i s i n keeping with the findings of Clark and Turner (7) to the effect that when KH^ PO^  solutions were allowed to react with solid calcium carbonate, the values of pH-%pCa were always reasonably close to the equilibrium value. Since there was a state of saturation with respect to calcite in the soils with the lower calcium carbonate content, i t would seem that where there was a small amount of CaCOg relative to s o i l , the s o i l influenced the solubility of the CaCG3 possibly by contaminating the surfaces of the CaC03 particles. On the other hand, when the CaC03 was present in larger amounts, there seemed to be enough CaC03 to overcome the influence of the s o i l , so that the f i n a l condition reflected equilibrium with CaC03 at the given partial pressure of carbon dioxide. 32 Table 13. The values of pH-%pCa and pH + pH 2P0 4 for 'calcareous' soils of varying GaC03 content at p c o 2 = 0.00011 atm. Calcium Carbonate Content % Sample pH @ 23.5°C pCa pH 2P0 4 pH-%pCa pH + pH 2P0 4 2 A 7.55 2.5453 3.3098 6.28 10.8598 B 7.62 2.6262 3.3054 6.31 10.9254 5 A 7.66 2.5818 3.3054 6.37 10.9654 B 7.66 2.5818 3.3401 6.37 11.0001 10 A 7.76 2.5817 3.3098 6.47 11.0698 B 7.76 2.5817 3.3098 6.47 11.0698 15 A 7.78 2.6128 3.2604 6.47 11.0404 Although the solutions were vir t u a l l y at equilibrium with solid calcium carbonate, they were not at equilibrium with a calcium phosphate compound. There was l i t t l e v a r i a b i l i t y in the phosphorus concentration in the percent of calcium carbonate originally present. Thus the value for pH2P04 was on the average equal to ,3.3. It i s seen from the diagram of Fig. 1 that under the conditions prevailing in this experiment, most of the points obtained by a plot of pH-%pCa and pH + pH2P04 clustered above the line for dicalcium phosphate. The inference i s that a state of supersaturation existed with respect to dicalcium phosphate, and that equilibrium was reached slowly i n these suspensions. The percent of calcium carbonate present i n the soils has l i t t l e effect on the equilibrium with respect to calcium phosphate compounds, provided that the volume of the dispersing solution and the Pco 2 w e r e constant. The indication i s that the reaction of phosphate with the s o i l was more important. 3 34 Part 3. ION RATIOS IN EXCHANGE RESIN SUSPENSIONS From the theory of the electric double layer i t was shown that the reduced ion ratios would remain constant throughout the double layer. If the concentration of ions at the surface of the exchanger does not change then the ion ratios would remain constant provided the electrolyte concentra-tion in solution is not markedly changed. Thus: (H*) = (H1") = constant /(Ca++) /(Ca++) before electrolyte after electrolyte addition.' addition. or pH-%pCa <= constant. As was discussed before,this relationship has been found to be constant in many so i l s . Since the indication is that there is not much exchangeable hydrogen in inorganic clays (1, 26, 33, 35), the constancy of the value of pH-%pCa in soils has been attributed to the proton transfer mechanism associated with hydrated aluminum ions. In soils i t i s d i f f i c u l t to determine whether this mechanism is operative since ions other than ct and if are l i k e l y to be present. To circumvent this d i f f i c u l t y a synthetic cation exchange resin was used instead of s o i l . The resin used was the weakly acidic carboxylic type, Amberlite IRC-50 (H). Large quantities of the resin were saturated with respect to calcium and hydrogen; washed free of electrolyte and then mixed i n quantities so as to give a predetermined percentage of saturation with respect to calcium. Cation exchange analyses were made on samples of the resin prepared i n this manner, and i t was found that the base saturation values were remarkably close to the calculated values. To weighed amounts of the calcium-saturated resin, CaCl2 was added in varying concentrations in a 1:2 resin-solution ratio. The samples were then shaken on a wrist action shaker for 18 hours, after which, pH was measured on the supernatant solution. Part of this solution was removed by f i l t r a t i o n and analysed for calcium. The relationship pH-%pCa was calculated and results are shown in the following tables. Table 14. The effect of increasing CaCl2 concentration on the value of pH-%pCa for Ca-H saturated resins. I n i t i a l Sample CaCl2 , Ca r m/LxlO-3 pH @ 23°C m/LxKT pCa pH-%pCa 2 U Ca-sat'd. IA 1.0 3.33 0.6748 3.2628 1.70 B 3.33 0.7024 3.2468 1.71 2A 1.5 3.23 1.1018 3.0703 1.70 B 3.24 1.0811 3.0670 1.71 3A 2.0 3.18 1.4839 2.9554 1.71 B 3.18 1.4839 2.9554 1.71 4A 2.5 3.15 1.9315 2.8549 1.72 B 3.14 1.9281 2.8557 1.71 5A 3.0 3.10 2.3412 2.7830 1.71 B 3.10 2.3412 2.7830 1.71 6A 3.5 3.09 2.7475 2.7283 1.73 B 3.08 2.7406 2.7242 1.72 "U i-sat'd. @ 22°C IA 1.0 3.46 0.9804 3.1113 1.90 B 3.48 0.7224 3.2333 1.86 2A 1.5 3.40 1.1696 3.0444 1.88 B 3.41 1.1730 3.0433 1.89 3A 2.0 3.37 1.6168 2.9200 1.89 B 3.36 1.6134 2.9210 1.90 )% t-sat'd @ 26°C IA 1.0 3.23 0.6818 3.2612 1.60 B 3.23 0.6818 3.2612 1.60 2A 1.5 3.15 1.0330 3.0979 1.60 B 3.15 1.0588 3.0881 1.61 3A 2.0 3.11 1.3894 2.9828 1.62 B 3.12 1.3637 2.9784 1.63 4A 2.5 3.06 1.8439 2.8749 1.62 B 3.06 1.8439 2.8749 1.62 5A 3.0 3.04 2.2261 2.8036 1.64 B 3.04 2.1951 2.8036 1.64 6A 3.5 3.02 2.6393 2.7401 1.65 B 3.02 2.6393 2.7401 1.65 Table 14 continued I n i t i a l CaCl, Ca Sample m/LxIO pH m/LxlO3 pCa pH-%pCa 65% Ca-sat'd @ 25° e 1A 1.0 3.33 0.7189 3.2380 1.71 B 3.33 0.6914 3.2536 1.70 2A 1.5 3.26 1.0836 3.0768 1.72 B 3.25 1.0836 3.0772 1.71 3A 2.0 3.19 1.4242 2.9707 1.71 B 3.19 1.4586 2.9615 1.71 4A 2.5 3.14 1.8851 2.8647 1.71 B 3.14 1.9298 2.8563 1.71 5A 3.0 3.12 2.2807 2.7933 1.72 B 3.12 2.2807 2.7933 1.72 6A 3.5 3.08 2.7279 2.7264 1.72 B 3.08 2.7279 2.7264 1.72 i% t-sat'd @ 23°C LA 1.0 3.39 0.7706 3.2088 1.79 B 3.39 0.7706 3.2088 1.79 2A 1.5 3.31 1.1421 3.0552 1.78 B 3.31 1.1367 3.0569 1.78 3A 2.0 3.26 1.5205 2.9454 1.79 B 3.26 1.5205 2.9454 1.79 4A 2.5 3.21 1.9746 2.8456 1.79 B 3.21 1.9746 2.8456 1.79 5A 3.0 3.18 2.3613 2.7786 1.79 B 3.18 2.3613 2.7786 1.79 6A 3.5 3.15 2.8139 2.7135 1.79 B 3.15 2.8139 2.7135 1.79 ;% i-sat'd. Q 25°C 1A 1.0 3.04 1.18 3.0481 1.52 B 3.04 1.18 3.0481 1.52 2A 1.5 2.82 1.61 2.9327 1.35 B 2.81 1.63 2.9284 1.35 3A 2.0 2.67 2.02 2.8506 1.25 B 2.67 2.02 2.8506 1.25 4A 2.5 2.57 2.49 2.7748 1.18 B 2.57 2.45 2.7804 1.18 5A 3.0 2.44 2.82 2.7333 1.07 B 2.42 2.65 2.7564 1.04 6A 3.5 2.34 3.15 2.6962 0.99 B 2.34 2.98 2.7178 0.98 37 In Table 14 the results are shown for the effect of increasing the CaCl2 concentration from 1.0 to 3.5mmoles/L on the value of pH-%pCa. In a l l cases, when the resins were saturated with from 21 to 85 percent calcium + (the other saturating ion being H), there was a decrease in pH with an increase in the i n i t i a l concentration of CaCl2 added. There was a slight decrease i n the i n i t i a l concentration of calcium after exchange had been effected, but the calcium concentration increased with a decrease in pH. The values for pH-%pCa remained constant over this range of saturation of the resin with respect to calcium. At 95% saturation with calcium, the pH decreased with an increase i n the i n i t i a l CaC^ concentration. The f i n a l calcium concentration increased with a decrease in pH. However, the value for pH-%pCa did not remain constant but changed 0.54 unit as the GaCl2 concentration was changed from 1.0 xlO" 3 m/L to 3.5 xl0"3m/L. In a similar experiment, the resin was saturated with calcium and aluminum, and pH and pCa determined as before. The concentration of aluminum was also determined on the f i l t r a t e s , by the Aluminon method (23). Details of the method are given in Appendix IV. The values for pH-%pCa are given in Table 15. Consideration was given taken that aluminum could interfere with the determination of calcium by the EDTA method with Cal-Red. However i t was found that the concentrations of aluminum in the f i l t r a t e s (maximum aluminum concentration = 0.0001 mmole/30ml.) would not interfere significantly with the calcium determination to cause error. 38 Table 15. The effect of increasing CaCl2 concentration on the value of pH-%pCa for Ca-Al saturated resins (47% saturated with Al.) I n i t i a l Ca Ca Al Sample m/LxlO3 pH @ 23°C m/Lxl03 pCa m/Lxl05 pH-%pCa 1A 1.0 3.63 0.8319 3.1762 0.4944 2.04 B 3.63 0.8319 3.1762 0.4944 2.04 2A 1.5 3.60 1.1899 3.0366 0.7416 2.08 B 3.60 1.1899 3.0366 0.7416 2.08 3A 2.0 3.57 1.6134 2.9200 0.7416 2.08 B 3.57 1.6134 2.9200 0.7416 2.11 4A 2.5 3.52 2.0168 2.8358 1.9775 2.11 B 3.52 2.0168 2.8358 1.9775 2.11 5A 3.0 3.45 2.4150 2.7689 2.2250 2.07 B 3.46 2.4150 2.7689 2.2250 2.08 6A 3.5 3.47 2.8134 2.7121 2.2250 2.11 B 3.46 3.1210 2.6784 2.2250 2.12 Table 16. Effect of Aluminum concentration on the determination Calcium by the EDTA Method with Cal-Red the indicator. Al Calcium taken. Calcium found. mmoles/30ml. mmoles. mmoles. % decrease. 0.0333 0.0999 0.1998 0.3333 1.0000 *** 0.025 0.025 0.025 0.025 0.025 0.0247 0.0247 0.0244 *** *** 1.20 1.20 1.20 At these concentration of aluminum, i t was impossible to record meaningful calcium concentrations. It was observed that the end-point of the titr a t i o n was reached when about 20% of the calcium present was titrated. However, the colour at the end-point did not persist but kept changing with time and subsequent additions of EDTA. 39 From Table 15 i t can be seen that the pH decreased with an increase in the CaC^ concentration. Calcium and aluminum concentrations increased with a decrease i n pH. The values for pH-%pCa were found to vary slightly with a maximum difference of 0.08. Within the limits of experimental error, particularly that of pH ( i t was d i f f i c u l t to measure the pH in these suspensions, possibly because of 'poisoning' of the electrodes by the aluminum), i t can be stated that the values for pH-%pCa were relatively constant. While this data in i t s e l f is not sufficient to establish beyond question that the constancy of the value of pH-%pCa can be attributed to the proton transfer mechanism associated with hydrated aluminum ions, the indication i s that there is a distinct possibility that this mechanism may be responsible for maintaining pH-%pCa constant in soils where there is a predominant amount of exchangeable aluminum. 40 SUMMARY In the study of the solubility of calcium carbonate in calcareous soi l s , i t was found that the calcareous materials behaved differently from calcite. These carbonate materials had fixed and unique solubilities that accounted for the constancy of the relationship pH-%pCa. Varying the s o i l -water ratios or the i n i t i a l CaC^ concentration did not affect the constancy of the values for pH-%pCa. The addition of increments of hydrogen saturated cation exchange resins and dilute HCl affected the constancy of the values for pH-%pCa. In the study of the fixation of phosphorus i n calcareous soils varying in calcium carbonate content, i t was found that the i n i t i a l fixation was rapid and that fixation continued long after the addition of phosphate. Although CaCOg exerted i t s normal effect on the solution equilibrium, no equilibrium was attained with respect to a calcium phosphate. Application of solubility c r i t e r i a indicated a state of supersaturation with respect to dicalcium phosphate. In resins saturated with calcium and hydrogen, the ion ratios remained constant when saturation was effected with calcium over the range 21 to 85%. At 95% calcium saturation the values for pH-%pCa were not constant. In a calcium-aluminum saturated resin the values for pH»%pCa were found to be relatively constant. APPENDIX I EQUATIONS FOR DEVELOPING SOLUBILITY DIAGRAMS FOR THE CALCIUM PHOSPHATES The lines for the calcium phosphates shown i n Fig. 1, were derived in the following manner; (Dicalcium phosphate dihydrate is used as an example) (1) pHP04 - pk 2 + pH 2P0 4 - pH . . . . . . Substituting (2) in (1) pCa + pk 2 + pH 2 P0 4 - pH = p k d c p d pCa + pH 2P0 4 = pH - p k d c p d - pk 2 -2(pH-%pCa) + pH + pH 2P0 4 - pk d c p d-pk 2 (2) (3) Evaluating (3) (pH + pH 2P0 4) » 2(pH-%pCa) + pk d c p d-pk 2 pH + pH 2P0 4 - 2(pH-%pCa) + 6.56 - 7.20 pH + pH 2 P0 4 = 2(pH-%pCa) - 0.64 . . . (4) APPENDIX II BUFFER SOLUTIONS FOR STANDARDISATION OF THE pH METER (51) M/20 Potassium Hydrogen Phthalate Solution pHt = 4.000 + h (t-15/100) 2 t being any temperature between 0°C. and 60°C. In d i s t i l l e d water i n a standard 1-litre volumetric flask dissolve 10.207 grams of potassium hydrogen phthalate A.R. grade, accurately weighed; make up to volume with d i s t i l l e d water and shake well. M/20 Sodium Borate Solution pHt » 9.27 - 0.0085 (t-15) t being any temperature between 0°C. and 60°C. In C0 2 free d i s t i l l e d water in a standard 1-litre volumetric flask dissolve 19.071 grams of Na2B407. 10 H20, A.R. grade, accurately weighed; make up to volume with C0 2 free d i s t i l l e d water and shake well. MOLYBDENUM-BLUE METHOD FOR THE DETERMINATION OF PHOSPHORUS. (16) Reagents, 1. Sulfomolybdic Acid Solution, 2.5%. Exactly 25.0 gm of c.p. ammonium molybdate, (NH4)6Mo7024.4 H20, is dissolved in 200 ml of d i s t i l l e d water and warmed to 60°C. Then 275 ml of phosphorus-free and arsenic-free concentrated sulfuric acid (35 to 36 N) is diluted to 750 ml with d i s t i l l e d water. After both solutions have cofcled, the ammonium molybdate solution i s - added slowly, with st i r r i n g , to the sulfuric acid solution. After the combined solution has cooled to room temperature, i t i s diluted with water to exactly 1000 ml. This i s a 9.7 to 9.9 N sulfuric acid solution, containing 2.5 gm of ammonium molybdate per 100 ml. 2. Chlorostannous acid Reductant. Approximately 25 gm of reagent grade SnCl^. 2H20 i s dissolved in 50 ml of concentrated HCl, with warming i f necessary to dissolve. This solution is diluted (with rapid stirring) to approximately 500 ml with recently boiled d i s t i l l e d water, giving about 0.2 M Sn. 3. Phosphorus Solution, 50 ppm. Dissolve 0.2195 gm of potassium dihydrogen phosphate RH2PO in about 400 ml of d i s t i l l e d water in a 1000-ml volumetric flask. Then 25 ml of 7 N HjSO^ i s added, and the solution made up to volume. Procedure. 1. An aliquot is placed in a 50-ml graduated test tube and the reaction adjusted to pH 3, using 2, 4-dinitrophenol indicator (0.25 percent in H 20). 2. Add 2ml sulfomolybdic solution. 3. The solution is made up to volume and 3 drops of the chlorostannous acid added. 4. The colour is allowed to develop for 4 minutes and then read photometrically at 660-mu. APPENDIX IV. 44 DETERMINATION OF CALCIUM BY EDTA TITRATION WITH CAL-RED. (32) Reagents. 1. Standard Calcium Chloride.0.01 N. Dissolve 0.5005 gram of pure dried CaC03 in 10 ml of approximately 3 N (1+3) HCl acid and dilute to a volume of 1 l i t r e . . 2. Potassium Cyanide. 3. Hydroxylamine Hydrochloride. 4. Potassium Hydoxide Solution. 8 N. 5. EDTA. (approx. 0.01N) Dissolve 2.00 gm. of disodium dihydrogen ethlenediaminetetra-acetate and 0.05 gm of magnesium chloride hexahydrate in water and dilute to 1 l i t e r . Standardize against Calcium Standard. 6. Cal-Red indicator. Procedure. 1. Pipetsuitable aliquot of solution containing calcium into a 125 ml. erlenmyer flask. 2. Add 4 ml. 8 N KOH. 3. Add about 30 milligrams KCN. 4. Add 30 milligrams Hydroxylamine Hydrochloride. 5. Add 50 milligrams Cal-Red indicator. 6. Titrate with EDTA to clear blue end-point APPENDIX V. METHOD EOR ALUMINUM DETERMINATION (23). 1. Pipeta 5 ml. aliquot of the f i l t r a t e into a test tube graduated at 25 ml. and add 2.5 ml. of 2.26 N HCl. 2. Add 5 drops of 1:20 thioglycollic acid (prepared daily by diluting 20 drops of 80% thioglycollic acid with 15 ml. of water). 3. Add 1 ml. 3% starch solution (prepared daily). 4. Add 5 ml. 3 N sodium acetate solution and s t i r . 5. Add 5 ml. 0.033 aluminon and s t i r . Make up the volume to 25 ml. and mix well. 6. Heat in a boiling water bath for exactly 4 minutes, remove, and cool in running water. Allow the tubes to stand for 30 minutes at room temperature. 7. Measure the transmittancy at 515 mu using 1 cm. c e l l and d i s t i l l e d water as blank. 8. Prepare a series of standards containing 0, 1, 3, 10, 20, 30, and 40 ugm. A l . (To prepare Al standards dissolve 0.2198 gm. uneffloresced A^SO^K^SO^,^ H20 in 1.13 N HCl and dilute to 500 ml. to give 25 ppm. Al in 1.13 N HCl. To prepare 1 ppm. and 10 ppm. Al make appropriate dilutions vith 1.13 N HCl). APPENDIX VI. 46 CALCULATION OF t^PO^ ACTIVITY. In a solution containing phosphate ions, the phosphate in solution can be represented as (Total Pj » CH3P04] + (H2PO4J + (HPO^ + fP0 4J (1) or (Total P} = (H 3P0 4) + (H 2P0 4) + (HPO5) + (Po | ) (2) cS where the brackets represent concentrations, the parentheses act i v i t i e s , and the symbol f is used for the activity coefficients. The thermodynamic dissocia-tion constants for phosphoric acid can be expressed as k x = (H*)(H 2P0 4) or (H 3P0 4) - (H +)(H 2P0 4) (3) (H3P04) k x k 2 - (H +)(HP0 4) or (HPOg) - (k 2)(H ?P0S) (4) (H 2P0 4) (H +) k3 - (H+XPOg) or (P0|) - kp^POj) (5) (HPO5) (H+)2 Substituting equations (3), (4), and (5) into equation (2) gibes (Total P] - (H +)(H 2P0 4) + (H 2P0 4) + k 2(H 2PO£) + k 2k 3(H 2P0 4) (6) k[ F (H*)f= (H+)z f= and solving for (H2P04) results i n (H2P04) » (Total P3 (7) (H +) + 1 + k 2 + k 2k3 kj_ f" (H+)f° (H+) 2f 2 APPENDIX VI continued, 47 CALCULATION 0FA1 + 3 ACTIVITY. In the determination of A l , the methods used gives the total concentra-tion in solution. As a result i t is necessary to know the dissociation constants of the various ion species in order to calculate Al activity. +3 .[4-In acid solutions the species Al and A10H~ predominate, but studies indicate that as the pH is raised a series of ions may form and eventually gels or hydrous aluminum oxides precipitate. (Brasset, 1952). In acid solutions, the total Al can be represented thus: (Total A l J - [A1 + 3J + fAlOH+^J (1) or (Total Al) • (Al+3) + (A10H**) " (2) f A l + 3 fAlOH"** Since the dissociation constant for the reaction A1(H 20)J 3 - A10H(H20)f" + H* (3) is known, then k^ can be expressed k x - (AlOHt1") (H +) or (AlOH^) » ki(Al+3) (4) (A1+3) (H+) Substituting equation (4) into equation (2) (Total A l J - (A1+3) + _ k 1 ( A l + 3 ) (5) f A l+3 (H^fMQH^ Solving for ( A l + 3 ) (A1+3) * (Total A l l (6) 1 + kx f l l * 3 (tfr)fA10H-H" Only the term 1 of the denominator is necessary below pH 3.2. *A1+3 48 BIBLIOGRAPHY 1. Aldrich, D.G. and Buchanan, J. R. 1958.. Anomalies in techniques for preparing H-Bentonites. S o i l S c i . Soc. Amer. Proc. 22:281-285. 2. Barbier, G. and Tyskiewicz, E. 1952. Mobilite des ions phosphoriques "fixes" dans le sol etudiee au moyen de P 2^' Int. Soc. S o i l S c i . Trans.2:79-83. 3. Brasset, C. 1952. On the reactions of the aluminum ion with water. Acta Chemica Scand. 6:910-940. 4. Brooks, R, Clark, L.M, and Thurston, E.F. 1950. Calcium Carbonate and i t s hydrates. Trans. Roy. Soc. Lon. A.243:145-167. 5. Chai Moo Cho and Caldwell, A.C. 1959. Forms of Phosphorus and Fixation in s o i l s . S o i l S c i . Amer. Proc. 23:458-460. 6. Clark, J.S. and Peech, M. 1955. Solubility c r i t e r i a for the existence of calcium and aluminum phosphates in s o i l s . S o i l S c i . Soc. Amer. Proc. 19: 171-174. 7. , and Turner, R.C. 1955. Reactions between solid calcium carbon-ate and orthophosphate solutions. Can. J. Chem. 33: 665-671. 8. Cole, C.V. 1957. Hydrogen and calcium relationships of calcareous s o i l s . S o i l S c i . 83: 141-150 9. Davis, F.L. 1943. Retention of phosphates by s o i l s . 1. Effect of addition of FeCl 3 and A l C l j upon the retention of phosphorus by virgin Hammond v . f . s . l . S o i l S c i . 56:457-478. 10. Ford, M.C. 1933. The nature of phosphate formation in soils J. Amer. Soc. Agron. 25: 134-143. 11. Frear, G. L. and Johnston, J. 1929. The solubility of calcium carbonate in certain aqueous solutions at 25°C. J. Amer. Chem. Soc. 51: 2082-2093. 12. Glasstone, 1946. S. Textbook of Physical Chemistry Edition 2. D. Van Nastrand Co., New York. 49 13. Harned, H.S. and Davis, R. 1943. The ionization constant of carbonic acid in water and the solubility of carbon dioxide in water and aqueous solution from 0 to 50°C. J. Araer. Chem. Soc. 65: 2030-2037. 14. , and Scholes, S. R. 1941. The ionization constant of HCO3 f r o m 0 t o 5 0 ° G ' J. Amer.Chem. Soc. 77: 271-279. 15. Haseman, H.F., Brown, E.H. and Whitt, CD. 1950. Some reactions of phosphate with clay and hydrous oxides of iron and aluminum So i l S c i . 70: 257-271. 16. Jackson, M.L. 1958. S o i l Chemical Analysis. Prentice-Hall Inc., Englewood C l i f f s . N.Y. 17. Kelly, J.B. and Mideley, A.R. 1943. Phosphate fixation - an exchange of phosphate and hydroxyl ions. S o i l S c i . 55: 167-176. 18. Kielland, J. 1937. Individual activity coefficients of ions in aqueous solutions. J. Amer. Chem. Soc. 59: 1675. 19. Klotz, L.M. 1950. Chemical Thermodynamics. Prentice-Hall Inc. Englewood C l i f f s . N.Y. 20. Kruyt, H.R. 1952. Colloid Science. Vol. 1. Elseview Pub. Co., N.Y. 21. Larsen, S. and Court, M.N. 1961. S o i l Phosphate solubility. Nature 189: 164-165. 22. Lekr, J.R. and Brown, W.E. 1958. Calcium phosphate f e r t i l i z e r s 11. A petro-graphic study of their alteration in s o i l . S o i l S c i . Soc. Amer. Proc. 22: 29-33 23. Lindsay, W.L. 1956 Ph.D. Thesis. Cornell University. 24. , and Moreno, E.C. 1960. Phosphate phase equilibria in s o i l s . S o i l S c i . Soc. Amer. Proc. 24: 176-182. 25. , Peech, M. and Clark, J.S. 1959. Solubility c r i t e r i a for the existence of variscite in s o i l s . S o i l S c i . Soc. Amer. Proc. 23: 357-360 26. Low, P.F. 1955. The role of aluminum in the ti t r a t i o n of bentonite. S o i l S c i . Soc. Amer. Proc. 19: 135-139 27. Mackenzie, J.E. 1923. Calcium carbonate hexahydrate. J. Chem. Soc. 123: 2409-2417. 28. Moser, F. 1942. Fixation and recovery of phosphate from some l a t e r i t i c s o i l s . S o i l S c i . Soc. Amer. Proc. 6: 328-334. 29. Nicol, W. E. and Turner, R. C. 1957. The pH of non-calcareous near neutral s o i l s . Can. J. S o i l S c i . 37: 96-101. 30. Olsen, S. K. and Watanabe, F.S. 1959. Solubility of CaCOo in calcareous s o i l s . S o i l S c i . 89: 288-291. 31. , and Cole, C.V. 1960. Solubility of phosphorus in calcareous s o i l s . S o i l S c i . 89: 288-291. 32. Patton, J. and Reeder, W. 1956. New indicator for t i t r a t i o n of calcium with (ethylenedinitrlo)-tetracetate. Anal. Chem. 28:1026-1028. 33. Russell, E. 1950. S o i l conditions and Plant Growth. Longmans, Green and Co., London. 34. Schofield, R.K. 1947. A ratio law governing the equilibrium of cations in solution. Proc. 11th. Int. Cong. Pure and Applied Chem. Lon. 3:25. 50 35. 1949. The effect of pH on electric charge carried by clay particles. J. So i l S c i . 1 :l-8 36. , and Taylor, A.W. 1955. The measurement of s o i l pH. So i l Sci. Soc. Amer. Proc. 19: 164-167. 37. Schollenberger, C.J. 1945. Determination of carbonate in s o i l . S o i l S c i . 59: 57-63. 38. Simmons, C F . 1939. The effect of carbon dioxide pressure upon the equilibrium system hydrogen colloidal clay -H2O - CaC03. J. Amer. Soc. Agron. 31: 638-648 39. Sveshnikova, V.N. 5 1 1952. Solubility of dolomitized limestone in H20 at 25°C. Doklady Akad.Nauk. U.S.S.R. 85: 357-360. 40. Terasvoori, A. 1930. Valtion Maatalouskoetorminnan. Julkaisuja No. 29. Helsinki. 41. Turner, R.C. 1958. A theoretical treatment of the pH of calcareous s o i l s . S o i l Sci. 86: 32-34. 42. , and Clark, J.S., 1956. The pH of calcareous s o i l s . S o i l S c i . 82: 337-341. 43. , and Nicol, W.E. 1958. The pH of strongly acid s o i l s . Can. J. S o i l S c i . 38: 63-68. 44. Whitney, R. S. and Gardner, R. 1943. The effect of carbon dioxide on s o i l reaction. S o i l S c i . 55:127-141. 45. Wild, A. 1954. The concentration of phosphate in the s o i l solution. Trans. F i f t h . Int. Cong. S o i l S c i . 2:500-504. 46. Wright, B. and Peech, M. 1960. Characterization of phosphate reactions products in acid soils by the application of solubility c r i t e r i a . S o i l S c i . 90: 32-43. 47. Yaalon, D. H. 1954. Physico-Chemical relationship of CaCO-j, pH and C0» in calcareous s o i l s . Trans. Intern. Congr. S o i l S c i . 5th. Congr. 2: 357-363. 48. 1957. Problems of s o i l testing on calcareous s o i l s . Plant and S o i l VII. 3: 275-288. 49. Yanat'eva, O.K. 1956. Application of the diagram of solubility of the system CaC03-MgC03-H2O for the characterization of certain carbonate rocks. J. Appl. Chem. USSR 29: 1247-53. 50. 1955. Solubility isotherms at 0° and 55°C. for the system Ca, Mg//C03S04-H20. Izvest. Sektora F i z . Khim. Anal., Akad Nauk. USSR 26: 266-9. 51. , Cambridge Instruction Booklet. 40. 52. , Saline and Al k a l i Soils. Handbook 60. U.S.D.A. 

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