CHEMICAL EQUILIBRIA

IN SOILS WILLARD L. L~DSAY Centennial Professor Colorado State University, Fort Collins

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CONTENTS

Symbols

1 Introduction

1.1 Dynamic Eqnilibria in Soils, 1.2 Intensity and Capacity Factors, 1.3 Elemental Composition of Soils,

References, Problems,

2 Methods of Handling Chemical Equilibria l 2.1 Equilibrium Constants, i

2,2 Concentration Versus Activity Cons\ants, 2.3 Ionic Strength, 2.4 Activity Coefficients, 2,5 Debye-Hiickel Equations, , 2.6 Activity Coefficients from Electrical :Conductivities, 2.7 Transforming Equilibrium Constant~, 2,8 Equilibrium Constants from Tbermqdynamic Data, 2.9 Redox Relationships, 2.10pe + pH, 2.11 E" versus lo.g K", 2.12 pe versus Eh,

xyii

1

2 4 6

8 9

10

11 12

. 12

13 13 16 18 21 23 23 26 27

ix

"

i i :

x CONTENTS CONTENTS xi

2.13 Redox Measurements in Natural Environments 28 7 Calcium 86

References, 30 7.1 CalCium Silicates and Alnrninosilicates, 87 Problems, 31 7.2 Other Calcium Minerals, 93

7.3 Complexes of Calcium in Solution, 95 3 Aluminum 34 7.4 Redox Relationships of Calcium, 98

3.1 Solubility of Alnminnm Oxides and Hydroxides, 35 7.5 The CaCO,-C02-H2 0 System, ,98 7.6 The Phase Rule, 101 3.2 Solubility of Aluminum Sulfates, 38 7.7 The CO2-H20 System, 102 3.3 Hydrolysis of AIH 39

3.4 Fluoride Complexes of Aluminum, 41 7.8 The CaO-C02-H20 System, 102

3.5 Other Aluminum Complexes, 43 7.9 The CaO-C02-H20-H2S04 System, 102

3.6 Estimating AI"" Activity, 45 Problems, 103 3.7 Redox Relationships of Alnminum, 46 3.8 Exchangeable Aluminum, 47 8 Magnesium 105

References, 48 8.1 Solubility of Magnesium Silicates, 106 Problems, 49 8.2 Magnesium Aluminosilicates, 112

4 Silica 50 8.3 Oxides,' Hydroxides, Carbonates, and Sulfates, 113 8.4 Magnesium Complexes in Solution, , 114

4.1 Forms of Silica in Soils, 51 8.5' Effect of Redox on Magnesium, 116 4.2 Silicate Species in Solution, 53

Problems, 116 References, 54 Problems, 55

9 Sodium and Potassium 118 5 Aluminosilicate Minerals 56 9.1 Solubility of Sodium Minerals, 119

5.1 Unsubstituted Aluminosilicates, 57 9.2 Solubility of Potassium Minerals, 123 5.2 Sodium Aluminosilicates, 62 9.3', Complexes of Sodium and Potassium, 125 5.3 Potassium Aluminosilicates, 65 9.4 : ReddxRelationships; 126 5.4 Calcium Aluminosilicates, 66

Problems, 5.5 Magnesium Aluminosilicates, 68 127

5.6 Summary Stability Diagrams for Aluminosilicates, 71 5.7 Controls of AIH Activity in Soils, 73 10 Iron 128 5.8 General Discussion of Aluminosi1icates, 75 10. I Solubility of Fe(IiI} Oxides in Soils, 129

References, 76 10.2 Other Fe(III} Minerals, 133 Problems, 77 10.3 Hydrolysis of Fe(III), 134

10.4 Fe(III) Complexes in Soils, . 136 6 Carbonate Equilibria 78 L 10.5 Effect of Redox on Fe(II)Sblubility, 139

6.1 The CO2-H2 0 'System, 79 I 10,6 Effect of Redox on the Stability of Iron Minerals, 141 6.2 The CO2-Soil System, 84 10.7 Hydrolysis and Complexes of Fe(II), 146

References, 84 i References, 148 Problems, . 84 ! ,Problems, 149

-------------~-~-,,~ •..... --.,~.~ ... -,.-

xii CONTENTS CONTENTS xiii

11 Manganese 150 15 Chelate Equilibria 238

. 11.1 Effect of Redox and pH on Manganese Solnbility, 151 15.1 Metal Chelates and Their Stability Constants, 239 11.2 Solution Species of Manganese, 157 15.2 Development of Stability-pH Diagrams for Chelates, 244

References, 160 15.3 Effect of Redox on Metal Chelate Stability, 252

Problems, 160 15.4 Chelation in Hydroponics, 256 15.5 Use of Chelates to Estimate Metal Ion Activities in Soils, 259

12 Phosphates 162 15.6 Use of Chelating Agents as Soil Tests, 261 15.7 Natural Chelates in Soils, 263

12.1 Orthophosphoric Acid, 163 12.2 Aluminum Phosphate~, 169 References, 264

12.3 Iron Phosphates, 173 ; Problems, 265

12.4 Effect of Redox on the Stability of Iron and Aluminum Phosphates, 177 16 Nitrogen 267

12.5 Solubility of Calcium Phosphates, 180 12.6 Effect of Redox on the Stability of Calcium Phosphates, 185 16. t Oxidation States of Nitrogen, 268

12.7 Solubility of Magnesium Phosphates, 186 ! 16.2 Equilibrium Betweeu Atmospheric N2 and O2, 269

12.8 Manganese Phosphates, 187 16.3 Effect of Redox on Nitrogen Stability, 272

12.9 Other Orthophosphates, 189 References, 279 12.10 Reduced Forms of Phosphorus, 189 ProbleII1B, 280 12.11 Stability of Polyphosphates in Soils, 190 12.12 Orthophosphate Complexes in Solution, 195 12.13 Reactions of Phosphate Fertilizers. with Soils,. f97 17 Sulfur 281

References, 204 I 17.1 Effect of Redox on Sulfur· Speciation, 282 Problems, 205 f 17.2 Dissociation of Sulfur Acids, 287

! 17.3 Formation of Elemental ~ulfur in Soils, 288

13 Zinc 210 17,4 Formation of Metal Sulfides, 290

1· 17.5 Effect of Sulfides on Meta,l Solubilities, 295 13.1 Oxidation State of Zinc 211 13.2 Solubility of Zinc Minerals in Soils, 211 References, 297 13.3 Zinc Species in Solution, 216 Problems, 297

References, 219 i Problems, 219

1

18 Silver 299

14 Copper 221 18.1 Effect pf Redox on the Stability of Silver Minerals, 300 14.1 Solubility of Cu(II) Minerals in Soils, 222 18.2 Solubility of Silver Halides and Sulfides, 304 14.1 Hydrolysis and Solution Complexes of Cu(Il), 228 18.3 Stability of Other Silver Miuerals, 306 14.3 Effect of Redox on Copper, 231 18.4 Stability of Silver Halide Complexes, 308 14,4 Complexes of Cu(I), 234 18.5 Hydrolysis Species and Other Silver Complexes, 310

References, 235 References, 313 Problems. 236 Problems, 313

___________ J __ _

xiv

19 Cadmium 19.1 Oxidation States of Cadmium in Soils, 19.2 Cadmium Minerals in Soils, 19.3 Hydrolysis Species of Cd(II), 19.4 Halide and Ammonia Complexes of Cadmium, 19.5 Other Cadmium Complexes, 19.6 Need for Further Studies,

References, Prohlems,

20 Lead

20.1 Solubility of Lead Minerals, 20.2 Hydrolysis Species of Lead, 20.3 Halide Complexes of Lead, 20.4 Other Complexes of Lead,

References, Problems,

21 Mercury

21.1 Stability of Hg(II) Minerals and Complexes, ~J.2 Stability of Hg(I) Minerals and Complexes, 21.3 Stability of Elemental Mercury, 21.4 Solubility of Mercury Sulfides in Soils, 21.5 Summary Redox Diagram for Mercury, 21.6 Organic Mercury Reactions,

References, Problems,

22 Molybdenum

22.1 Molybdenum Species in Solution, 22.2 Stability of Molybdenum Minerals in Soils, 22.3 The Elfe,t of Redox on Molybdenum Solubility,

References, Problems,

23 Organic Transformations

23.1 Oxidation States of Carbon; 23.2 Products of Glucose Metabolism,

CONTENTS

315 316 316 321 322 323 326

326 327

328

329 338 339 341

341 342

343

344 353 355 358 359 362

362 362

364

365 367 369

372 372

373

. 374 375

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CONTENTS

23.3 Reactions of Acetic Acid, 23.4 . Oxidation to CO2 (g) and Reduction to CH4(g); 23.5 Stability of Graphite,

References, Problems,

Appendix Standard Free Energies of Formation

Index

x

38( 38: 38:

38: 38:

38,

42,

w TABLE 20.1 EQUlLIBRIUM REACTIONS FOR LEAD MlNERALSAND COMPLEXES AT 25"C w 0

Reaction No. Equilibrium Reaction log KC

Oxides, Carbonates, Sulfates

1 PbO(yel!ow) + 2H+ ¢ PbH + H,O 12.89 2 PbO(red) + 2H+ ¢ Pb2+ + H,O 12.72 3 Pb(OH),(c) + 2H+ ¢ PbH + 2H,O 8.16 4 Pb;O.(c) + 8H+ + 2e- ¢3Pb2+ + 4H,O 73.79 5 PbO,(c) + 4H+ + 2e- ¢PbH + 2H,O 49.68 6 PbC03(cerussite) + 2H+ ¢ PbH + CO,(g) + H,O 4.65 7 Pb,C03CI,(p'bosgenite) + 2H+ ¢ 2Pb2+ + CO,(g) + H,O + 2CI- -1.80 8 Pb3(C03),(OH),(c) + 6H+ ¢3Pb2+ + 2CO,(g) + 4H,O 17.51 9 PbC03 ·PbO(c) + 4H+ ¢2PbH + CO,(g) + 2H,O 17.39.

10 PbSO.(anglesite)¢Pb2+ + SO~- -7.79 11 PbSO.·PbO(c) + 2H+ ¢2Pb2+ + SO~- + H,O -0.19 12 PbSO.·2PbO(c) + 4H+ ¢3Pb2+ + SO~- + 2H,O 11.01 13 PbSO.·JPbO + 6W¢4Pb2+ + SOi- + 3H,O 22.30

Silicates

14 PbSi03(c) + 2H+ + H,O ¢ Pb2+ + H,SiO~ 5.94 15 Pb,SiO.(c) + 4H+ ¢ 2Pb2+ + H.SiOr 18.45

Phosphates

'·f •

. ".... ,-.,.,_ ... ","'"".,-- ," " ... -.. ~~.-.--.....---' .... -.......... ---"'--... -.. , ... 23"": + 6H+¢5 +F- -12.98

Other Minerals

24 soil.Pb ¢ Pb2+ -8.50' 25 PbMoO.(wulfenite)¢PbH + MoO~- -16.04 26 PbS(galena) ¢ Pb" + S'- -27.51 27 PbH + 2e- ¢ Pb(c) -4.33

H~drol~sis Species

28 PbH + H,O¢PbOH+ + H+ -7.70 I 29 PbH + 2H,O¢Pb(OH), + 2H+ -17.75 I

,

30 PbH + 3H,O¢Pb(OH)3 + 3H+ -28.09 " 31 PbH + 4H,O¢Pb(OH)~- + 4H+ -39.49

32 2PbH + H,O¢Pb,OHH + H+ -6.40 33 3Pb2+ + 4H,O¢Pb3(OH)~+ + 4H+ -23.89 34 4Pb2+ + 4H,O ¢ Pb.(OH)!+ + 4H+ -20.89 35 6Pb'+ + 8H,O¢Pb6(OH)~+ + 8H+ -43.58

Halide Complexes

36 Pb2+ + Br- ¢ PbBr+ 1.77 w 37 Pb2 + + 2Br- ¢ PbBr~ 2.60 ;:;

38 Pb2+ + 3Br- ¢ PbBr, 3.00 (Continued)

340 CH.20 LEAD

this development was soil-Pb at 10- 8 •5 [14 Pb2+. In general the halide complexes have similar stabilities but decrease slightly in the orderPbl + > PbBr+ > PbCI - > PbP-. At halide activities > 10-4 M, all of these complexes begin to contribute significantly to total lead in solution. Iu

- the halide range of 10- 2 to 10-1.5 111 the halide complexes are approximately equal to free PbH

. Only in the halide activity range > 1O~ 1.3 M do the higher order complexes such as PbI3, PbI" Pbli-, PbF" PbBr~, aud PbCI~ exceed the simple 1: 1 complexes. Generally the halide complexes. of Pb2+ are not highly significant in soils. '

Although Fig. 20.4 was developed for a'Pb2+ activity of 10- 8 .5 111, it can be used to obtain halide complexes at any Pb2+ activity. Increasing or

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-8

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Pb 2 +

log 5°-42 -.,

log ~042--,

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log NOg ., .

',. log NO) -31 PbCS04)2 -, ~

·PbH2P04t 1 -3

'0 .f!

~'o/

., '4 / G ... _4 Pb( N03)2

4 5 6 pH

7 8 9

Fig. 20.5 Phosphate, nitrate, nnd sulfate complexes of lead in equilibrium with soil-Ph and PbCO.,(cerussite) at 0.003 alm of CO2 _ The PbP10?- line is ill metastable equilibrium with. fl-CU;!P20~ (el tlnd soil-Ca.

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334 cn.20 LEAD

-4

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, N .0 0.. -8

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-9

I ·K Kaolinite

1- G Gibbsfte -10 Q Quartz

5 Soil-Si

-11

-12 4 5 6 7 8 9 10

pH

Fig.20.2 The solubility of various lead silicates llnd phosphates compared to PbC03tcerlissite) when phosphute is controlled by various solid phases as. indicuted and CO:.!.(g) is 0.003 lHm.

in Fig. 20.2. Nriagn (1974) examined the solubilities of many lead phosphates and pointed out how they may be important in controlling lead in natural environments. Since the average lead content of soils (10 ppm) is generally much less than tliat of phosphorus (600 ppm), it is highly/possible th~t phosphate may control PbH solubility.

In developing Fig. 20.2 phosphate activity was fixed by strengite and soil-Fe at low pH, by P-tricaJcium phosphate (P-TCP) and soil-Ca at intermediate pH, and by P-TCP, calcite, and 0.003 atm of COz at high pH (Fig. 12.8). The P-TCP reference was arbitrarily selected as intermediate between the phosphate maintained by CaHP04 • 2HzO(DCPD) as an llpper limit and by hydroxyapatite(HAP) as a lower limit. Tlw dashed lines in Fig. 20.2 show the effect of these limits on PbH activity as phosphate shifts equilib­rium from DCPD to HA.

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I=QUUUU NM"'<t+"''''''" "'<t NM"<t+ +++ ++++++ ++++ +;!; + + + ':b~;!;;!; ';b;!;;!;;!; 1'1 ,.0 M <'I M P-t.D.o.D P-t ,.0 .;? .0 .DP-t.o,.o,.o P-tP-t~ ~_P-t

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SOLUBILITY 01' LEAD MINERALS 333

-3

-4

'PbSOd (cH-'3

-5

., -6

• • .<J £L -7 OJ

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-Se- I C02 (g~

-9 e- a 0.003 tm b 0.OQ03 atm

-10

-11 --, 4 5 6 7 S 9 10 pH

Fig.20.1 The solubility of various lead oxides, carbonates, and sulfates when SOi- and Cl­are 10-3 M and CO2 is 0.003 atm or as specified.

Oflthe1minerals included in Fig. 20.1 PbSO,(anglesite) is the most stable below pH 6,f,Whereas PbC03(cerussite) is most ~table at higher pH values. Changing slopes and shifts in the sOlnbility lines for the various minerals in Fig. 20.1 are logical reflections of the ratios df Pb/SO" Pb/Cl, Pb/C0

3 present iq these minerals.

The solubilities of several lead silicates and -phosphates are given by Reactions 14 throngh 23 of Table 20.I-and are plqtled in Fig. 20.2. Cernssite is inc1ud<;d for comparison. The silicates Pb2 SiO,(c) and PbSi0

3(c) lines

are shownJor equilibrium with quartz at 10-4 M H4SiO~ aud with shifts indicated:for soil-Si and kaolinite-gibbsite (Chapter 5). These lead silicates are too soluble to be expected in soils.

Lead is capable of forming numerous phosphate minerals as shown

TABLE 19.1 EQUILIBRIUM CONSTANTS OF CADMIUM MINERALS AND COMPLEXES .

Reaction No.

1

2 3 4 5 6 7 8 9

10 11 12

13 14 15 16 17 18 19 20

21 22 23 24 25 26 27 28 29

318

Equilibrium Reaction

Redox Reaction

Cd2+ + 2e-~ O\(c)

Minerals

CdO(monteponite) + 2H+ ~CdH + H,O P-Cd(OH),(c) + 2H+ ~ Cd2+ + 2H20

CdCO,(oct.vite) + 2H+ ~ CdH + CO2 (g) + H2 0 CdSiO,(c) + 2H+ + H20 ~ CdH + H4SiO'

CdS04(c)~Cd2+ + SO;:­CdS04.H20(C)~CdH + SO;:- + H20

CdS04.2Cd(OH),(c) + 4H+ ~3Cd2+ + SO::- + 4H20

2CdS04

.Cd(0H)'(c) + 2H+ ~3Cd2+ + 2S01- + 2H,O Cd,(P0

4),(c) + 4H+ ¢3Cd2+ + 2H,PO'

CdS(greennokite)¢ CdH + S2-soil-Cd ¢ CdH

Hydrolysis Species

Cd2+ + H20 ¢ CdOH+ + H+ CdH + 2H,O¢ Cd(OH)1 + 2H+ CdH + 3H,O¢Cd(OH), + 3H+ Cd2+ + 4H,O¢Cd(OH);,- + 4H+ Cd2+ + 5H,O ¢ Cd(oH)l- + 5H+ CdH + 6H,O¢Cd(OH)~- + 6H+ 2Cd'+ + H,O¢Cd,OH3+ + H+

4Cd2+ + 4H,O ¢ Cd4 (0H)1+ + 4H+

Halide Complexes

Cd2+ + Br-¢CdBr+ Cd2+ + 2Br- ¢ CdBr, Cd2+ + 3 Br- ¢ CdBr, Cd2+ + 4Br- ¢ CdBrl-

CdH + Cl- ¢CdCI+ Cd2+"+ 2Cl- ¢CdCl, Cd2+ + 3Cl- ¢CdCI, Cd2+ + 4CI- ¢CdCll,-

CdH + r- ¢ Cdr+

logKO

-13.64

15.14 13.65 6.16 7.63

-0.04 -1.59 22.65

6.73 1.00

-27.D7 -7.00*

-10.10 -20.30 -33.01 -47.29 -61.93 -76;81 -6.40

-27.n

2.15 3.00 3.00 2.90 1.98 2.60 2.40 2.50 2.28

! I,

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I

CADMIUM MINERALS IN SOILS

TABLE 19.1 (Colltinued)

Reaction No.

30 31 32

33 34 35 36

37 38 39 40 41 42 43

Equilibrium Reaction

Cd2 + + 21- ¢ CdI2 Cd" + 3 r- ¢ CdI, Cd2+ + 4r- ¢ CdIi-

Ammonia Complexes

. Cd2 + + NHt :;;=-CdNH5+ + H+ Cd2+ + 2NHt ¢ Cd(NH,)j+ + 2H+ Cd2+ + 3NHt ¢Cd(NH,)j+ + 3H+ Cd" + 4NHt ¢ Cd(NH,)i+ + 4H+

Other Co.mplexes

CdH + CO,(g) + H,O¢CdHCOj + H+ Cd" + C02(g) + H,O¢CdCO; + 2H+

Cd2+ + NO, ¢ CdNOj CdH + 2NO,¢Cd(NO,),

Cd2+ + H2PO' ¢ CdHPO' + H+ Gd2+ + P2oi- ¢CdP20~-

Cd2+ + SO;,- ¢ CdSO'

• Estimated from Street et al. (1978).

319

logK'

3.92 5.00 6.00

-6.73 -14.00 -21.95 -30.39

-5.73 -14.06

0.31 0.00

-4.00 8:70 2.45

Also included iu Fig. 19.1 is Cd3(PO.h(c) which is plotted for five different conditions depending on which minerals control phosphate. The selected phosphate controls include: FePO.· 2H20(strengite) and Fe(OHh(soil) at low pH, !l-Ca3(PO.)2(tricalcinm phosphate) and soil-Ca at intermediate pH, and tricalcium phosphate (TCP), CaC03(calcite), and CO2(g) at high pH. In acid soils Cd3(PO,)2(c) is too soluble to account for the Cd2+ levels generally found. In the pH range of 6 to 7.5, Cd,(PO,),(c) in eqnilibrium with TCP controls Cd2+ at approximately 10-5

.5 M. 10 calcareous soils octavite is

more stable than Cd,(p04),(c) in eqUilibrium with TCP or more insoluble phosphate minerals.

In recently fertilized soils CaHPO.· 2H20(brushite) may be present and support a higher level of phosphate than does TCP (Chapter 12). Under these conditions, Cd,(PO.J,(c) increases in stability as shown by the dashed lines in Fig. 19.1, and Cd2+ would be depressed to approximately 10- 6 to 10- 7 M, in the pH range of 6.3 to 7.4. Above this pH, Cd3(PO.J,(c)

--------------------~-------------------------------------------------------

324 CH. 19 CADMIUM

-6

-7

-8

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OJ .Q

-11

c12

log 'Stol~ . I I ~2 1 CdC0 3 --.I

-.SOil-Cd----T------r---toctavite) I Cd 2+ I I ,

CdSO.:1° 1-3

rog Noi .,

_3 Cd N03+

'·4 . 0"'· ,p

c.;0 .

0"

log 1'103-. ,

{J

-131 ,/Cd(N03l2 \., \

-14 I I I \ )1

4 5 6 7 8 9 10 pH

Fig. 19.4 Cadmium complexes of carbonate, nitrate, phosphate, and sulfate in eql.\ilibrium with soil-Cd and CdCO;t(octavite) at 0,003 atm CO 2 .

Equation 19.11 is plotted as the CdHCOj line in Fig. 19.4. When Cd2+ is controlled by octavite .

CdC03(octavite) + 2H+ Cd2+ + CO2 (g) + H 2 0

~ ~ Cd2+ + CO2(g) + H20

CdHCOj + H+

CdCO.(octavite) + H+ =" CdHCOj

10gKO

6.16 -5.73

0.43

(19.12)

: CADMIUM MINERALS IN SOILS

~ '[j

-2

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-4

-5

u. -6 OJ .Q

\ \ . ""',r.oo

-7 I Soil-Cd

'..f'~.J*" '\

-9 ..

-9

a C02 '" O.OO~ atm b. cb2 = 0·0003 atrfl '* Ca 2+" 10- 2

•5 M .

, \ \-\ r\ \

-10' \ 4 5 6 7

pH 9 9

317

10

Fig. "19.t The sol~bil~ty of several cadmium minerals cO,mpared to soil-Cd at 10- 1 M.

· included here was estimated from Street et aI. (1978), who fouud Cd2+ activities of approximately 10- 7 M iu the pH range of 6 to 7.5. Their findings are sUIIl!i:mrized b~ the reaction j

soil-Cd . Cd2+ log KO = -7.00 (19.3)

, At pH values above 7.5; depending on CO,(g), Cd2+ activity is limited by CdC03(octavite). As shown in Fig. 19.1, octavite with a CO2 of 0.003 atm i dep,esses Cd2+ 100-fold for each unit increase in pH. At this C02 Ie,'el, the 1 depression begins at pH 7.84. · The mineral CdSi03(c) is more soluble than octavite, so it is not expected to form in soils. The Ihinerals CdO(monteponite), p·Cd(OHh(c) and'

· CdS04 · 2Cd(OH)2(c) are also too soluble to form in soils.

TABLE 14.1 EQU1LIBRIUM CONSTANTS FOR VARIOUS REACTIONS IN­VOLVING COPPER

Reaction No.

2 3 4 5

6 7

8 9

\0 11 12

13 14 15

16 17 18 19 20

21 22 23

224

Equilibrium Reaction

Cu(II) Oxides, Hydroxides, and Carbonates

CuO(tenorite) + 2H+:::=: Cu2 + + H 2 0 Cu(OHh(c) + 2H+ ¢ CU" + 2H,O

CuCO,(C) + 2H+ ¢ Cu" + COz(g) + H,O CU2(0H)zCO,(malachite) + 4H+ ¢2Cu2+ + CO2(g) + 3H20 \ Cu,(OH),(CO,h(azurite) + 6H+ ¢3Cu" + 2C02(g) + 4H20\

CU(lI) Ferrite and Soil-Cli

a-CuFe204(cupric ferrite) + 8H+ ¢Cu2 + + 2Fe3 + + 4HzO Soil-Cu_+ 2H+ ¢ Cu2 +

, Cu(II) Sulfates

CilS04(eha1coeyanite)¢Cu" + SO~-CuSO •. 5HzO(e)¢Cu" + SOl- + 5HzO

CuO·CuSO.(e) + 2H+ ¢3Cu2+ + SOl- + H,,o Cu4(OH)6S0.(bronehantite) + 6H+ ¢4Cu" + SOl- + 6H20

CU.(OH)6S04·1.3H,O(c) + 6H+ ¢4Cu" + S01- + 7.3H20

Cu(Il) Phosphates

Cu,(p04lz(e) + 4H+ ¢3Cu" + 2H2PO;;: Cu,(PO.)z·2HzO(e) + 4H+ ¢3Cu2+ + 2H2PO;;: + 2H,0

CU2P207(C)¢2Cu2+ + P20~-

CU(lI) Hydrolysis

CUB + H,O¢CuOH+ + H+ CUB + 2H,O ¢ CU(0H)2 + 2H+ Cu" + 3H,0¢Cu(OH), + 3H+ Cu" + 4H~0¢Cu(OH)i- + 4H+'

2Cu" + 2H20¢CU2(OH)~+ + 2H+

Cu(I!) Complexes

Cu2 + + Cl- ¢ CuCl+ Cu" + 2Cl- ¢ CuCI~ Cu2+.+ 3CI-¢CuCl,

)ogKO

7.66 8.68 8.52

12.99 19.57

10.13 2.80

3.72 "-2.61 11.50 ' 15.35 17.27

2.24 0.34

-15.22

-7.70 -13.78 -26.75 -39.59 -10.68

0.40 -0.12 -1.57

~! i' I,

~ 1,\ '~I '1,

f' !: .,' I ,t, ;f

{

1\, ,1; ~,

i~~: f/·

\i ··'i!

SOLUBILITY OF Co(II) MINERALS IN SOlLS

TABLE 14.1

Reaction No.

24 25 26 27 28 29 30 31 32 33 34 35

36 37 38

39 40 41 42

43 44 45 46

,

(Continued)

Equilibrium Reaction

Cu" + CO2(g) + H,O ¢ CuHCOj + H+ Cu" + CO2(g) + H20 ¢ CuCO~ + 2H+

Cu2+ + 2CO,(g) + 2H,0¢Cu(CO,)l- + 4H+ Cu2 + + NO; ~ CuNOi'

. Cu~+ + 2NO)" ~ CU(N0 3)2 Cu2 + + H 2POi ¢·CuH2PO; Cu2 + + H2 POi:;;::: CuHP04 + H+

Cu2 + + 2H+ + P20j- ¢ CUH2P20~ Cu2 + + H+ + P20j- :;;:::CUHP207

Cu2 + + P20~-¢CUP20~,-2Cu 2 +P;zOj- ¢-CU2P2 0 7 Cu2 + + S01- ¢ CUS04

Redox Reactions

Cu2 + + e- ¢ Cu+ Cu+ + e- ¢Cu(c)

Cu" + 2e- ¢ Cu(e)

Cu(I) Minerals

CU20(euperite) + 2H+ ¢2Cu+ + H20 CuOH(e) + H+ ¢Cu+ + H,O

IX-Cu2Fe204(cuprous ferrite) + 8H+' ¢ 2Cu+ + 2Fe~+ + 4H20 Cu2S04(e)¢29u+ + SOl'

Cu(I) Complexes

Cu+ +. Cl- ¢ CuCio

Cu+ + 2Cl-¢CuCI, Cu+ + 3CI-¢CuClj-

2Cu+ + 4CI- ¢Cu,CI~-

,

225

logKO

-5.73 -11.43 -26.48

0.50 -0.40

1.59 -4.00 18.67 14.78 6.64

-0.03 2.36

2.62 8.87

11.49

-2.17 -0.70

-13.53 -1.95

2.70 5.51 5.70

13.10

than phosphorus (Table 1.1), this plot was (ieve10ped to show how phos­phorus may affeci the solubility ofCu", The mineral CU3(P04h ·2I;t,O(c) is always more stable than Cn3(PO.),(C) as it lowers Cu" by O.6Jlog unit. Defails of only the dihydrate mineral are shown in Fig. 14.3. In this plot phosphate is fixed by strengite and soil-Fe at low pH, by hydroxyapatite, (3-TCP, or DCPD with soil-Ca at intermediate pH, and by the same minerals with calcite and CO,(g) at high pH.

D

<t' 0::'

',>

'>/ / \.()

// 0 '-~/ !1:~'V

nO~"'// nOf)'V \: / ~ ....

V",/ 0"'''> G/'

o N , '<t , <D ro , , <I:znJ 501

(,v .;::.."

",0

Q

(,v .~' .

",0'

E E ~~

.m", . . .1J.l .; N lL. '0 '0 - ~ ~ a NN

00 • uu ..... n:I t I "51u 1IJ IV .c 1 ::!:! +' ~:=..!:::!~ +' 0 III rII U)tnUU rn.cu-o

N 'T

'<t 'T

II' / I o

226

N , '<t , <D ro ,

• cn:) 50,

o , N 'T :! ,

en ~ @ B

ro I I ~ N ..s::

"" i I ..8 ;0. IJ.o

J, ,

oIl

'<t

en

ro

I'-

<D

oIl

'<t

• "0

~ j ~ "l ;!;

(i!

" ~ i ~

1l <l!

~ ·0

I I U,d 0. ni U

~~. .... 0 o ~

,£ '" ~ ,S

~ "I ;!;

~

, \

SOLUBILiTY OF Cu(II) MINERALS IN SOILS 227

Both copper phosphates are more soluble than soil-Cu especially as phosphate solubility is depressed by hydroxyapatite. These copper phosphates can dissolve sufficiently in soils to provide both available copper and phosphorns for plants.

The solnbility of Cn2P207(C) is given by Reaction 15, and the stabilities of several copper pyrophosphate complexes are given by Reactions 31 through 34 of Table 14.1. These stahility relationships are plotted in Fig. 14.4 under the conditions that pyrophosphate is fixed by !3-Ca2P207(C) and soil-Ca '(Section 12.10) and Cn2+ is fixed by soil-Cu. The complexes CuHP20" CUH2P20~, CUP20~-, and CU2P207 are all less abundant than the H2P20~-, C.aHP20:;-, and CaP20~- shown in the upper part of

,., :!: .eo +'

o

-2

-4

u -6

'" !J1 .Q

-10

-12

-14.

,­P20 7

"...- CU2P207°

/ I

, I

./-f-P20;-(CU2P207-soil Cu)

/

,/caH P207- ,\. CaP20

-16 I \, \, "" ~ 4 5 6 7 8

pH

Fig.14.4 The insignificance of copper pyrophosphate compIexe~ in soils when [J-Ca2P20 7(C) and soit-Ca fix P20i:-. The instability OfCU2P20,(C) relative to /1-Ca1PJO,{c) is also shown .

i !

____ .L.....----___ .

230 CH. 14 COPPER

-6

-8

.0" . '.

-10 ~;;,"

"''' '~'7 C'~"k:.

cu(oH)i

" ~'i " ~O~ 'tOg C02 ,CUC03'

~~. ~ 0 -~,t;:>O" -12

Q ,0 '0- ~-S: "1 !.

-, .. ~.~ ,.,

!.I 0 ~a ~

~,).~ C:l.t

.;;: :c;

~ °0_ u -14

\.~ .

oJ

O~~ ~ m

G , c> \"

.Q

.. .' ~ , . ~ ~

-16

, . ~'11. v C}. ",,0

d' . -18 r-:.

6

~ ",",,/

Q. '. -20 f-

:G ~. -22 1 I

4 5 6 7 8 9 pH

Fig.14.6 Various complexes ofOl2+ in equilibrium with soil-Cu.

1 % to total copper in solution include CuRCO;" at high CO2 and neutral pH, CuHPO' at high phosphate and neutral pR, and Cu(CO,)l- at high CO2 and high pH Complexes such as CuNO;", CU(N03)2, CuH2POt, Cu(CO,)i-, CuCI+, CuCI", and CuCI, do not contribute significantly to total copper at the level of anions normally found in soils.

Total inorganic copper in soil solution [Cu;.".] can be fairly well rep­resented by the equation

[Cu;.".] = [Cu2+] + [CuSO'] + [Cu(OH),] + [CuCO~] (14.5)

Substituting activities gives

(Cu2+) [Cu;.".] = --+ (cuso') + (CU(OH)2) + (CuC03) (14.6)

'YCu2+

EFFECT OF REDOX ON COPPER 231

When each term on the-right side of this equation is expressed in terms of its equilibrium constant given in Table 14.1, Eq. 14.6 can be rearranged to give:

[Cu;.".]. 10 -;'---::-:;:~~:-~1~013. 78 + (Cu2+) = _1 + 102.'6(SO~ ) + (H+r YCuH

11.43C0 2(g) (H+)'

(14.7)

Thus the activity of Cu2 + can be e~timated from this equation with measure-" ments of (1) total inorganic copper in solution, (2) pH, (3) ionic strength, (4) the activity of SO~-, and (5) the partial pressure of CO2 (g). If conditions are such that other inorganic copper complexes contribute significantly to soluble copper, they must also be included in Eq. 14.5 and reflected in Eq.14.7.

Copper is normally Present in soil solution as both organic and inorganic complexes (Hodgson et al., 1966). For this reason total soluble copper is expressed as

[Total soluble copper] = [Cu;.",] + [Cn".] (14.8)

. and some independent means must be used to distinguish between the two forms of copper. As indicated in Section 13.2, soils can be extracted in the presence and absence of carbon black which absorbs many of the soluble organic constituents. Intl1is way Cu2+ activities can be estimated and used to test the solubility rehitionships depicted in Fig. 14.1.

14.3 EFFECT OF REDOX ON COPPER

Copper equilibria which are affected by redox are given by Reactions 36 through 38 of Table 14.1. According to Reaction 36:

Cu2+ + e'" ==== Cu+ log KO = 2.62 (14.9)

Cu+ . log Cu2+ = 2.62 - pe (14.10)

Thus Cu + and Cu2+ activities are equal at pe = 2.62, and their ratio changes lO-fold for each unit change in pe.

What are the stable copper minerals in soils which become reduced? Figure 14.7 was developed to answer this question. Soil-Cu defined by Reaction 14.1 is stable above pe + pH of 14.89. Below this redox Cu2Fe,04 (cuprous ferrite) in eqnilibria with soil-Fe becomes the stable copper min;ral.

214 CH.13 ZINC

0

" -2 ~~~r~~' 'O~ -J-

"

... \. ....

", -4

-6 • N

C N

'" .Q -8

-10

-12

---14 -.3 4 5 6 7

pH

8 9 10

Fig. 13.1 The solubilities of several zinc minerals compared to soil-Zn.

increase in pH. The effect of different levels ofH4SiO' (controlled by quartz, soil-Si, or amorphous silica) and of CO,(g) on the solubilities of these zinc minerals is also indicated.

Earlier Lindsay and Norvell (1969) hypothesized that ZnSiO,(c) in equililJ­rium with SiO,(amorp) may accountfor the Zn2+ solubility in soils depicted by Eq. 13.4. Further investigations into the solnbility of ZnSi03 (c) reported in the literatnre showed that value to be in error (Norvell and Lindsay, 1970). From Fig. 13.1 it appears that ZnFe,O.(franklinite) may account for the solubility of Zn2+ depicted by soil-Zn. The soil-Zn line can be expected to shift somewhat for different soils. Likewise the solubility of franklinite shifts depending ou the activity ofFe3+. For example, Fe(OHh(amorp) depresses the solubility of franklinite, whereas crystalline Fe(III) oxides such as maghemite or goethite lower Fe3+ and permit higher equilibrium' levels of Zn2+ in Fig. 13.1. Added to this complexity is the fact that franklinite may be amorphous and other cations of similar size may partially substitute for

I

r ,

1.

I I'

;~ i. I,

.~ I f

I'

I

SOLUBILITY OF ZINC MINERALS IN SOILS 215

Zn2+. Thus it is easy to see why Zn2+ solubility in soils can result in a range orsolubilities. Iron oxides form analogous minerals with other trace elements like Cu2 +, Mn_2-1', Co 2 +, etc. Further investigations are needed to examine the more. general applicability of the soil-Zn level of soluble Zn2+ in soils, and the possibility that franklinite in equilibrium with Fe(III) oxides

. possibly control Zn2+ solubility in soils. All of the Zn(OH). minerals, ZnO(zincite), and ZnC03 (smithsonite)

shown in Fig. 13.1 are too soluble to persist in soils. They are about 10' times more soluble than soii-Zn. These minerals make good zinc fertilizers in soils becanse they dissolve sufficiently to' maintain levels of Zn2+ that are adequate for plants.· The mineral Zn2Si04(wiIlemite) is of intermediate solubility, but it is too soluble to account for the soil-Zn found in most soils.

,

The solubilities of various zinc chlorides, sulfates, oxysulfates, and phos­phates are given by Reactions 13 through 17 of Table 13.1. The log KO values of these minerals indicate that greater than 1 M levels of Zn2+ would be necessary for. ZnCI,(g), ZnS04(zinkosite); and ZnO· 2ZnS04(c) to form in soils. Hence these minerals are not expected'in soils.

, " c N

o

-2

-4

'" -6 .Q

-8

-10

"' 0 .... ;-

'", 'Q

a Strengite - Soil Fe b Soil-Ca c Calcite (17) Zn3(P04)204H20

b

o

-12' '\!

4 5 6 7 8 9 pH

Fig. 13.2 The solubilities of hopeite and ZnO . ZnS04 (c) compared to other zinc minerals when phosphate is fixed by various iron and calcium _phosphates (drawn for CO2{g) =

0.0003 .tm).

218

.P :~

-2

-4

-6

-8

tJ -10 rn ill Q

-12

-14

-16

CH,13 ZINC

~&

o/\~Oi;,.:z~.,

<1> '. .~~" ,'", '\,', "

O'>~ f": I~ r ' "'f' '~ 0" 4''<{ " 'g:, ~ <1>,. . ~

'''r; ~ , <',,>. ~; ... ' "OL, • 1,., 'x'" '7

' 0 " <o'tb •

"

6-,OC?> •

~ C} ,9--<',

:.

. ~ .

~h{' .

op'>i ' ,; ~

; C}~

~' "

,cf? ',J.

. <'''" •

6-,<1> '

~ , .0:,

-18 I I I . !

3456789 pH

Fig. 13.4 Chloride, nitrate, phosphate, and sulfute [;omplexes of ZnH in eqllilibrium with soil~Zn. .

ZINC SPECIES IN SOLUTION

.P :~

° -2

-4

-6

1J . -8 rn

'" ,Q

-10

-12

, -14

~i'~

</J~f.y'f. .

Zn,(OH)/

I ~ 4 5 6

~ ,-,-0 0''' 1." 0 ' ",'"

I IY ( . ' 7 8 9 10

pH

i Fig. 13.3 The hydrpiysis species of Zn2~ in equilibrium with soil-Zn.

. ~;"V".''''': ._"

~"'-.... ~~"-

-""~~!"" ... ","~~,- ... -

.... ,,--'

"

217