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- Herausgeber: Wiley-VCH
- Kategorie: Wissenschaft und neue Technologien
- Sprache: Englisch
- Veröffentlichungsjahr: 2016

Endlich ein Fachbuch, das die Theorie, Methoden und die verschiedenen Arten von Metall-Ionen-Komplexen in Wasser (Hydrolyse) umfassend behandelt. Geschrieben wurde dieses Referenzwerk von einem Kernchemiker aus dem Hochschulbereich und einem Geochemiker aus der Industrie. Behandelt werden Kationen- und Anionen-Komplexe sowie die Metall-Ionen-Hydrolyse, zu der zunächst Hintergrundinformationen geliefert werden, bevor eine Beschreibung der Dissoziation von Wasser, aller verschiedenen Hydrolysekomplexe und Verbindungen von Metall und Wasser folgt. Ein Muss für Wissenschaftler im universitären Umfeld und in der Industrie, die sich mit diesem interdisziplinären Thema beschäftigen.

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Cover

Title Page

Copyright

Preface

Volume 1

1 Introduction

References

2 Theory

2.1 Hydrolysis Reactions and Stability/Solubility Constants

2.2 Debye–Hückel Theory

2.3 Osmotic Coefficient

2.4 Specific Ion Interaction Theory

2.5 Determination of Temperature-Dependent Parameters

2.6 Determination of Ion Interaction Parameters from Activity and Osmotic Coefficient Data

2.7 Determination of Ion Interaction Parameters for KOH at Temperatures Other than 25 °C

2.8 Activity of Water

2.9 Enthalpy and Entropy

2.10 Estimation of Stability and Solubility Constants

References

3 Methodologies for Determining Stability/Solubility Constants

3.1 Introduction

3.2 Potentiometry

3.3 Liquid–Liquid Extraction

3.4 Solid–Liquid Extraction

3.5 Solubility

3.6 Spectrophotometry

3.7 Experimental Uncertainties

References

4 Statistical Analysis and Selection Criteria

4.1 Uncertainty Assessment

4.2 Selection of Stability and Solubility Constants

References

5 Water

5.1 Physical Properties of Water

5.2 Protolysis of Water at Zero Ionic Strength

5.3 Protolysis of Water in Media of Fixed Ionic Strength at 25 °C

5.4 Protolysis of Water in Media of Fixed Ionic Strength at Other Temperatures

5.5 Enthalpy, Entropy and Heat Capacity

5.6 Collation and Assessment of Literature Data

5.7 Zero Ionic Strength Data at High Pressure

5.8 Comparative Strength of the Protolysis Constant of Water

References

6 Alkali Metals

6.1 Introduction

6.2 Lithium

6.3 Sodium

6.4 Potassium

6.5 Rubidium, Cesium and Francium

References

7 Alkaline Earth Metals

7.1 Beryllium

7.2 Magnesium

7.3 Calcium

7.4 Strontium

7.5 Barium

7.6 Radium

References

8 Scandium, Yttrium and the Lanthanide Metals

8.1 Scandium

8.2 Yttrium

8.3 Trivalent Lanthanide Metals

8.4 Cerium(IV)

References

9 Actinide Metals

9.1 Actinium

9.2 Protactinium

9.3 Uranium

9.4 Neptunium

9.5 Plutonium

9.6 Americium

9.7 Curium

9.8 Berkelium and Californium

References

Volume 2

10 Titanium(IV), Zirconium, Hafnium and Thorium

10.1 Titanium(IV)

10.2 Zirconium

10.3 Hafnium

10.4 Thorium

References

11 First Transition Series Metals

11.1 Titanium(III)

11.2 Vanadium

11.3 Chromium

11.4 Manganese

11.5 Iron

11.6 Cobalt

11.7 Nickel

11.8 Copper

11.9 Zinc

References

12 Second and Third Series Transition Metals

12.1 Introduction

12.2 Molybdenum

12.3 Technetium

12.4 Rhodium

12.5 Palladium

12.6 Silver

12.7 Cadmium

12.8 Iridium

12.9 Platinum

12.10 Gold

12.11 Mercury

References

13 Aluminium, Gallium, Indium and Thallium

13.1 Aluminium

13.2 Gallium

13.3 Indium

13.4 Thallium

References

14 Tin and Lead

14.1 Introduction

14.2 Tin

14.3 Lead

References

15 Bismuth and Polonium

15.1 Introduction

15.2 Bismuth

15.3 Polonium

References

16 Prediction of Stability and Solubility Constants

16.1 Theory

16.2 Prediction of Hydrolysis Stability Constants

16.3 Prediction of Solubility Constants of Oxide and Hydroxide Phases

16.4 Prediction of Stability Constants for Monomeric Species of Higher Stoichiometry

16.5 Prediction of Constants at Temperatures other than 25 °C

16.6 Application of the UTMIC in Assessment of Stability/Solubility Constant Data

References

Appendix: Extended Specific Ion Interaction Theory: Ion Interaction Coefficients

Index

End User License Agreement

Volume 1

2 Theory

Figure 2.1 Derivation of ion interaction parameters

ε

n

(K

+

, OH

−

) from (a) activity and (b) osmotic coefficients of KOH solutions. The solid line is the line of best fit, and the dotted lines are the 95% uncertainties projected out from

I

= 0–10 mol kg

−1

.

Figure 2.2 Derivation of ion interaction parameters

ε

n

(Na

+

, OH

−

) from (a) activity and (b) osmotic coefficients of NaOH solutions. The solid line is the line of best fit, and the dotted lines are the 95% uncertainties projected out from

I

= 0–10 mol kg

−1

.

3 Methodologies for Determining Stability/Solubility Constants

Figure 3.1 The hydrogen concentration of a solution in absence of metal (solid squares and solid line) and the hydrogen ion concentration in the presence of metal (open circles and dotted line) plotted against the amount of titrant added.

Figure 3.2 The difference between the free hydrogen ion concentration in a titrated solution without metal and that containing metal as a function of titrant added.

Figure 3.3 The average ligand number as a function of pH for three thallium(III) concentrations.

Figure 3.4 The average ligand number as a function of pH for four cerium(IV) concentrations. (The data are from Danesi (1966).) The differing symbols represent the data obtained at differing metal ion concentrations.

Figure 3.5 Extraction system using an acidic extractant.

Figure 3.6 Extraction system using a solvating ligand.

Figure 3.7 A typical extraction curve as a function of pH where an acidic extractant is utilised.

Figure 3.8 A typical sorption curve.

Figure 3.9 A typical solubility curve (solid line – data (solid squares) from Ekberg

et al.

(2004)). Also shown (dashed lines) are concentrations of individual species determined from the hydrolysis and solubility constant data given by Ekberg

et al.

(2004).

Figure 3.10 Spectra of the hydrated copper(II) ion and the Cu(NH

3

)

4

2+

complex.

Figure 3.11 Changes in the absorption spectra with increasing hydrolysis of the UO

2

2+

ion.

Figure 3.12 A fish bone structure for illustrating contributing uncertainties in the determination of a

D

ratio.

Figure 3.13 Comparison of independent and dependent variable space: regions 1 and 2 are independent, while region 3 is dependent.

Figure 3.14 A covariance matrix.

5 Water

Figure 5.1 The density of water as a function of temperature.

Figure 5.2 Boiling points of some ‘water lookalikes’.

Figure 5.3 Critical point (CP), boiling point (BP) and melting point (MP) temperatures of some of the second period elements.

Figure 5.4 Phase diagram of water.

Figure 5.5 Pourbaix diagram for water at 25 °C.

Figure 5.6 Enthalpy of the protolysis of water. The open circles are the data of Sweeton, Mesmer and Baes (1974) and the closed triangles those of Palmer and Drummond (1988). The line is determined from the protolysis constant data of Bandura and L'vov (2006).

Figure 5.7 Heat capacity of the protolysis of water. The open circles are the data of Sweeton, Mesmer and Baes (1974) and the closed triangles those of Palmer and Drummond (1988). The line is determined from the protolysis constant data of Bandura and L'vov (2006).

Figure 5.8 Comparison of average of measured data at each temperature (with respective 95% confidence interval) from Table 5.7 and predicted data from the equations of Bandura and L'vov (2006) (solid line).

Figure 5.9 Water activity data for cesium chloride media derived from the osmotic coefficient data of Robinson and Stokes (1959). The solid line is derived from use of Eq. (5.19).

Figure 5.10 Dependence of log

K

w

on ionic strength in cesium chloride media. The solid line is derived from use of Eq. (5.17).

Figure 5.11 Water activity data for TMA chloride media derived from the osmotic coefficient data of Lindenbaum and Boyd (1964). The solid line is derived from use of Eq. (5.19).

Figure 5.12 Dependence of log

K

w

on ionic strength in TMA chloride media. The solid line relates to Eq. (5.17).

Figure 5.13 Comparison of

a

1

values (Table 5.11) with those predicted from the use of Eq. (5.20).

Figure 5.14 Comparison of

a

2

values (Table 5.11) with those predicted from the use of Eq. (5.21).

Figure 5.15 Dependence of log

K

w

on ionic strength in sodium chloride media at 25 °C. The solid line relates to Eq. (5.17).

Figure 5.16 Comparison of

a

1

values (Table 5.16) with those predicted from the use of Eq. (5.20).

Figure 5.17 Comparison of

a

2

values (Table 5.16) with those predicted from the use of Eq. (5.21).

Figure 5.18 Dependence of log

K

w

on ionic strength in potassium chloride media at 25 °C. The solid line relates to Eq. (5.17).

Figure 5.19 Water activity data for lithium chloride media derived from the osmotic coefficient data of Robinson and Stokes (1959). The solid line is derived from use of Eq. (5.19).

Figure 5.20 Dependence of log

K

w

on ionic strength in lithium chloride media. The solid line is derived from use of Eq. (5.17).

Figure 5.21 Dependence of log

K

w

on ionic strength in barium chloride media at 25 °C. The solid line relates to Eq. (5.17).

Figure 5.22 Water activity data for sodium perchlorate media (reported in Brown, Curti and Grambow (2005) (solid squares)) and derived from osmotic coefficient data of Robinson and Stokes (1959) (open circles). The solid line is derived from use of Eq. (5.19).

Figure 5.23 Dependence of log

K

w

on ionic strength in sodium perchlorate media. The solid line relates to Eq. (5.17).

Figure 5.24 Water activity data for lithium perchlorate media derived from the osmotic coefficient data of Robinson and Stokes (1959). The solid line is derived from use of Eq. (5.19).

Figure 5.25 Dependence of log

K

w

on ionic strength in lithium perchlorate media. The solid line is derived from use of Eq. (5.17).

Figure 5.26 Dependence of log

K

w

on ionic strength in potassium iodide media. The solid line relates to Eq. (5.17).

Figure 5.27 Dependence of log

K

w

on ionic strength in potassium nitrate media. The solid line relates to Eq. (5.17).

Figure 5.28 Dependence of log

K

w

on ionic strength in sodium nitrate media. The solid line relates to Eq. (5.17).

Figure 5.29 Dependence of log

K

w

on ionic strength in sodium triflate media. The solid line relates to Eq. (5.17).

Figure 5.30 Dependence of log

K

w

on ionic strength in sodium sulphate media. The solid line relates to Eq. (5.17).

6 Alkali Metals

Figure 6.1 Stability constant (log *

β

1

°) of LiOH(aq) as a function of the reciprocal of absolute temperature.

Figure 6.2 Stability constant (log *

β

1

°) of NaOH(aq) as a function of the reciprocal of absolute temperature.

7 Alkaline Earth Metals

Figure 7.1 Structures of the beryllium hydrolysis species. The open bonds are in the plane of the page, the closed bonds come out of the page, and the hatched bonds go into the page. Each coloured area denotes a separate species.

Figure 7.2 Solubility constant (log *

K

s10

°) of α-Be(OH)

2

(s), β-Be(OH)

2

(s) and BeO(s) as a function of the reciprocal of absolute temperature.

Figure 7.3 Stability constant (log *

β

1

°) of BeOH

+

as a function of the reciprocal of absolute temperature.

Figure 7.4 Stability constant (log *

β

2

°) of Be(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 7.5 Stability constant (log *

β

3

°) of Be(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 7.6 Dependence of log *

β

2

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 7.7 Dependence of log *

β

21

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 7.8 Dependence of log *

β

33

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 7.9 Predominance diagram for the speciation of the beryllium(II) ion at 25 °C.

Figure 7.10 Predominance diagram for the speciation of the beryllium(II) ion at 100 °C.

Figure 7.11 Solubility constant (log *

K

s10

°) of Mg(OH)

2

(s) as a function of the reciprocal of absolute temperature.

Figure 7.12 Stability constant (log *

β

1

°) of MgOH

+

as a function of the reciprocal of absolute temperature.

Figure 7.13 Predominance diagram for the speciation of the magnesium(II) ion.

Figure 7.14 Solubility constant (log *

K

s10

°) of Ca(OH)

2

(s) as a function of the reciprocal of absolute temperature.

Figure 7.15 Stability constant (log *

β

1

°) of CaOH

+

as a function of the reciprocal of absolute temperature.

Figure 7.16 Measured solubility data of Ca(OH)

2

(s) in various ionic media from Johnston and Grove (1931) compared with calculated data using the extended specific ion interaction theory.

Figure 7.17 Measured solubility data of Ca(OH)

2

(s) in sodium nitrate media from Yeatts and Marshall (1967) compared with calculated data using the extended specific ion interaction theory.

Figure 7.18 Measured solubility data in NaOH and KOH media from Duchesne and Reardon (1995) compared with calculated data using the extended specific ion interaction theory.

Figure 7.19 Stability constant (log *

β

1

°) of SrOH

+

as a function of the reciprocal of absolute temperature.

Figure 7.20 Stability constant (log *

β

1

°) of BaOH

+

as a function of the reciprocal of absolute temperature.

8 Scandium, Yttrium and the Lanthanide Metals

Figure 8.1 Dependence of log *

β

1

of ScOH

2+

Figure 8.2 Dependence of log *

β

22

of Sc

2

(OH)

2

4+

Figure 8.3 Dependence of log *

β

35

of Sc

3

(OH)

5

4+

Figure 8.4 Predominance diagram for the speciation of the scandium(III) ion at 25 °C.

Figure 8.5 Speciation diagram of scandium (10

−4

mol kg

−1

) in the −log [H

+

] range of 3.0–5.2, using stability constants relating to (a) zero ionic strength and (b) 1.0 mol kg

−1

NaClO

4

.

Figure 8.6 Dependence of log *

β

1

of YOH

2+

Figure 8.7 Relationship between the solubility constant (log *

K

s10

°) of Ln(OH)

3

(s) and that of the stability constant of LnOH

2+

(log *

β

1

°). The data are given in Tables 8.11 and 8.12. The solid line of the relationship is a quadratic.

Figure 8.8 Solubility constant (log *

K

s10

°) of Ln(OH)

3

(s) (lanthanum, neodymium and gadolinium) as a function of the reciprocal of absolute temperature.

Figure 8.9 Stability constants (log *

β

1

° at 25 °C) for LnOH

2+

illustrating a clear tetrad effect (solid lines) and application of the refined spin-pairing energy theory to describe the constants (dashed line). Only three of the predicted stability constants lie outside the uncertainty limits of the selected log *

β

1

° values listed in Table 8.12. The ionic radii for the

x

-axis data come from Shannon (1976).

Figure 8.10 Dependence of log *

β

1

of NdOH

2+

Figure 8.11 Linear free energy relationship between the stability constants of the Ln

2

(OH)

2

4+

species and those of the LnOH

2+

species.

Figure 8.12 Dependence of log *

β

1

of LaOH

2+

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). The data in the open squares are the corrected data from other temperatures that clearly are not consistent with the other data shown in the figure.

Figure 8.13 Dependence of log *

β

1

of CeOH

2+

Figure 8.14 Dependence of log *

β

1

of PrOH

2+

Figure 8.15 Dependence of log *

β

1

of SmOH

2+

Figure 8.16 Dependence of log *

β

1

of EuOH

2+

Figure 8.17 Dependence of log *

β

1

of GdOH

2+

Figure 8.18 Dependence of log *

β

1

of TbOH

2+

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficient and stability constant at zero ionic strength).

Figure 8.19 Dependence of log *

β

1

of DyOH

2+

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficient and stability constant at zero ionic strength).

Figure 8.20 Dependence of log *

β

1

of HoOH

2+

Figure 8.21 Dependence of log *

β

1

of ErOH

2+

on ionic strength in sodium perchlorate media (the solid line is obtained using the derived interaction coefficient and stability constant at zero ionic strength).

Figure 8.22 Dependence of log *

β

1

of TmOH

2+

Figure 8.23 Dependence of log *

β

1

of YbOH

2+

Figure 8.24 Dependence of log *

β

1

of LuOH

2+

Figure 8.25 Predominance diagram for the speciation of the neodymium(III) ion at 25 °C.

Figure 8.26 Stability constant (log *

β

1

) of CeOH

3+

as a function of the reciprocal of absolute temperature from the studies of Hardwick and Robertson (1951) (open circles and dashed line; 2.20 mol l

−1

HClO

4

and 5–35 °C) and Everett and Skoog (1971) (solid squares and solid line; 1.05 mol l

−1

HClO

4

and 7.6–30.8 °C).

Figure 8.27 Stepwise stability constant (log *

K

2

) of Ce(OH)

2

2+

as a function of the reciprocal of absolute temperature from the study of Everett and Skoog (1971) in 1.05 mol kg

−1

HClO

4

.

Figure 8.28 Dimerisation stability constant (log *

K

D

) of Ce

2

(OH)

2

6+

as a function of the reciprocal of absolute temperature from the study of Hardwick and Robertson (1951) in 2.20 mol kg

−1

HClO

4

.

Figure 8.29 Dependence of log *

K

2

of Ce(OH)

2

2+

on ionic strength in perchloric acid media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 8.30 Dependence of log

K

D

relating to the formation of Ce

2

(OH)

2

6+

as a function of ionic strength in perchloric acid media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

9 Actinide Metals

Figure 9.1 Stability constant (log *

K

2

) for the formation of PaO(OH)

2

+

as a function of the reciprocal of absolute temperature.

Figure 9.2 The stability constant (log *

K

3

) for the formation of PaO(OH)

3

(aq) as a function of the reciprocal of absolute temperature. The open circle datum is from Trubert, Le Naour and Jaussaud (2002) (see text).

Figure 9.3 Dependence of log *

K

2

of PaO(OH)

2

+

on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.4 Dependence of log *

K

2

of PaO(OH)

3

(aq) on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). The data indicated as open circles are from Trubert, Le Naour and Jaussaud (2002, 2003) and Fourest

et al.

(2004) (see text).

Figure 9.5 The stability constant (log *

β

1

°) for the formation of UOH

3+

as a function of the reciprocal of absolute temperature.

Figure 9.6 Dependence of log *

β

1

of UOH

3+

on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.7 Dependence of log *

β

1

of UOH

3+

on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.8 The stability constant (log *

β

1

°) for the formation of UO

2

OH

+

as a function of the reciprocal of absolute temperature.

Figure 9.9 The stability constant (log *

β

22

°) for the formation of (UO

2

)

2

(OH)

2

2+

as a function of the reciprocal of absolute temperature.

Figure 9.10 The stability constant (log *

β

35

°) for the formation of (UO

2

)

3

(OH)

5

+

as a function of the reciprocal of absolute temperature.

Figure 9.11 Dependence of log *

β

1

of UO

2

OH

+

(aq) on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.12 Dependence of log *

β

1

of UO

2

OH

+

(aq) on ionic strength in nitrate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.13 Dependence of log *

β

22

of (UO

2

)

2

(OH)

2

2+

on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.14 Dependence of log *

β

22

of (UO

2

)

2

(OH)

2

2+

on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.15 Dependence of log *

β

22

of (UO

2

)

2

(OH)

2

2+

on ionic strength in nitrate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.16 Dependence of log *

β

34

of (UO

2

)

3

(OH)

4

2+

on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.17 Dependence of log *

β

34

of (UO

2

)

3

(OH)

4

2+

on ionic strength in nitrate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.18 Dependence of log *

β

35

of (UO

2

)

3

(OH)

5

+

Figure 9.19 Dependence of log *

β

35

of (UO

2

)

3

(OH)

5

+

Figure 9.20 Dependence of log *

β

35

of (UO

2

)

3

(OH)

5

+

on ionic strength in nitrate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.21 Predominance diagram for the speciation of the uranyl(VI) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 9.22 Dependence of log *

β

1

of NpOH

3+

Figure 9.23 Dependence of log *

K

s10

of NpO

2

OH(am) on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and solubility constant at zero ionic strength).

Figure 9.24 Dependence of log *

β

1

for NpO

2

OH(aq) on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.25 Dependence of log *

β

1

for NpO

2

(OH)

2

−

Figure 9.26 Dependence of log *

K

s10

of PuO

2

·H

2

O on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.27 Dependence of log *

β

1

of PuOH

3+

Figure 9.28 Dependence of log *

β

1

of AmOH

2+

on ionic strength in sodium chloride and sodium perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength for NaCl and the dotted line for NaClO

4

where Δ

ε

2

= 0).

Figure 9.29 Dependence of log *

β

2

of Am(OH)

2

+

on ionic strength in sodium chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.30 Dependence of log *

β

3

of Am(OH)

3

(aq) on ionic strength in sodium chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.31 Dependence of log

β

1

of CmOH

2+

on ionic strength in sodium chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Figure 9.32 Dependence of log

β

2

of Cm(OH)

2

+

on ionic strength in sodium chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength).

Volume 2

10 Titanium(IV), Zirconium, Hafnium and Thorium

Figure 10.1 Solubility data of TiO(OH)

2

(am) from the studies of Babko, Gridchina and Nabivanets (1962), Liberti, Chiantella and Corigliano (1963) and Nabivanets and Lukachina (1964).

Figure 10.2 Solubility constant (log *

K

s11

) for the formation of TiOOH

+

as a function of the reciprocal of absolute temperature.

Figure 10.3 Solubility constant (log

K

s12

) for the formation of TiO(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 10.4 Stability constant (log *

K

3

) for the stepwise formation of TiO(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 10.5 Predominance diagram for the speciation of the titanyl(II) ion at 25 °C. The behaviour in the region of 2 > −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 10.6 Dependence of log *

β

1

of ZrOH

3+

Figure 10.7 Dependence of log *

β

1

of ZrOH

3+

Figure 10.8 Dependence of log *

β

2

of Zr(OH)

2

2+

Figure 10.9 Dependence of log *

β

2

of Zr(OH)

2

2+

Figure 10.10 Dependence of log *

β

4

of Zr(OH)

4

(aq) on ionic strength in nitrate (solid squares: the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength) and perchlorate (open circles: the dashed line is obtained using the zero ionic strength stability constant derived from the analysis of the nitrate data and the derived interaction coefficients for perchlorate) media.

Figure 10.11 Dependence of log *

β

34

of Zr

3

(OH)

4

8+

on ionic strength in perchlorate media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). The dotted line is in relation to the selected stability constant at zero ionic strength.

Figure 10.12 Dependence of log *

β

34

of Zr

3

(OH)

4

8+

on ionic strength in chloride media (the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). The dotted line is in relation to the selected stability constant at zero ionic strength.

Figure 10.13 Fraction of zirconium removed onto a KU-2 cation exchange resin as a function of acid concentration and that predicted using Eq. (10.1) (solid line).

Figure 10.14 Comparison of measured and modelled (solid line) distribution ratio of zirconium from the study of Nazarenko and Mandzhgaladze (1969) as a function of hydroxide concentration from measurements in (a) 0.3 mol l

−1

(H,Na)ClO

4

and (b) 0.5 mol l

−1

(H,Na)ClO

4

.

Figure 10.15 Predominance diagram for the speciation of the zirconium(IV) ion at 25 °C and in 1.0 mol l

−1

(H,Na)ClO

4

. The behaviour in the region of 0 > −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 10.16 Stability constant (log *

β

48

) of Th

4

(OH)

8

8+

as a function of the reciprocal of absolute temperature.

Figure 10.17 Dependence of log *

β

1

of ThOH

3+

on ionic strength in perchlorate media (solid squares: the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). Also given are data from nitrate media (open circles).

Figure 10.18 Dependence of log *

β

22

of Th

2

(OH)

2

6+

on ionic strength in nitrate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.19 Dependence of log *

β

22

of Th

2

(OH)

2

6+

on ionic strength in chloride media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.20 Dependence of log *

β

22

of Th

2

(OH)

2

6+

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength. Δlog *

β

22

= log *

β

22

− log *

β

22

° and log *

β

22

° is the average of that obtained from nitrate and chloride media.

Figure 10.21 Dependence of log *

β

23

of Th

2

(OH)

3

5+

on ionic strength in chloride media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.22 Dependence of log *

β

48

of Th

4

(OH)

8

8+

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.23 Dependence of log *

β

6,15

of Th

6

(OH)

15

9+

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.24 Dependence of log *

β

6,15

of Th

6

(OH)

15

9+

on ionic strength in nitrate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 10.25 Predominance diagram for the speciation of thorium(IV) at 25 °C and zero ionic strength.

Figure 10.26 Predominance diagram for the speciation of the thorium(IV) ion at 25 °C and in 1.0 mol l

−1

NaClO

4

.

11 First Transition Series Metals

Figure 11.1 Dependence of log*

β

1

of TiOH

2+

on ionic strength in chloride media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.2 Dependence of log*

β

22

of Ti

2

(OH)

2

4+

Figure 11.3 Dependence of log*

β

1

of VOH

2+

Figure 11.4 Dependence of log*

β

2

of V(OH)

2

+

Figure 11.5 Dependence of log*

β

22

of V

2

(OH)

2

4+

Figure 11.6 Dependence of log*

β

1

of VOOH

+

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.7 Dependence of log*

β

22

of (VO)

2

(OH)

2

2+

Figure 11.8 Dependence of log*

β

1

of VO

2

OH(aq) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.9 Dependence of log*

β

2

of VO

2

(OH)

2

−

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength. Data obtained in chloride media are identified as open circles.

Figure 11.10 Dependence of log

K

3

of VO

2

(OH)

3

2−

on ionic strength in sodium media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.11 Dependence of log*

K

4

of VO

2

(OH)

4

3−

on ionic strength in sodium media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.12 Dependence of log*

K

D

for the dimerisation reaction and formation of (VO

2

)

2

(OH)

6

4−

as a function of ionic strength in sodium media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.13 Dependence of log*

K

D

for the dimerisation reaction and formation of (VO

2

)

2

(OH)

4

2−

on ionic strength in sodium media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.14 Dependence of log*

K

for the dimerisation reaction and formation of (VO

2

)

2

(OH)

5

3−

Figure 11.15 Dependence of log*

K

for the formation of (VO

2

)

4

(OH)

8

4−

Figure 11.16 Dependence of log*

K

for the formation of (VO

2

)

4

(OH)

9

5−

Figure 11.17 Dependence of log*

K

for the formation of (VO

2

)

4

(OH)

10

6−

Figure 11.18 Dependence of log*

K

for the formation of (VO

2

)

5

(OH)

10

5−

Figure 11.19 Dependence of log*

K

for the formation of (VO

2

)

10

(OH)

14

4−

on ionic strength in NaClO

4

media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.20 Dependence of log*

K

for the formation of (VO

2

)

10

(OH)

15

5−

Figure 11.21 Dependence of log*

K

for the formation of (VO

2

)

10

(OH)

16

6−

Figure 11.22 Predominance diagram for the speciation of the vanadium(V) ion at 25 °C. The behaviour in the regions, 2 > −log [H

+

] > 12, should be treated with caution due to changes in activity coefficients.

Figure 11.23 Predominance diagram for the speciation of the vanadium(V) ion at 25 °C and in 1.0 mol l

−1

NaClO

4

. The behaviour in the regions, 2 > −log [H

+

] > 12, should be treated with caution due to changes in activity coefficients.

Figure 11.24 Dependence of log*

β

1

of CrOH

2+

Figure 11.25 Dependence of log*

β

22

of Cr

2

(OH)

2

4+

Figure 11.26 Dependence of log*

β

34

of Cr

3

(OH)

4

5+

Figure 11.27 Stability constant (log*

β

1

) for CrOH

2+

as a function of the reciprocal of absolute temperature. Data are for zero ionic strength (), 0.068 mol l

−1

NaClO

4

(), 0.232 mol l

−1

NaClO

4

(×) and 0.966 mol l

−1

NaClO

4

().

Figure 11.28 Stability constant (log*

β

22

) for Cr

2

(OH)

2

4+

as a function of the reciprocal of absolute temperature.

Figure 11.29 Stability constant (log*

β

34

) for Cr

3

(OH)

4

5+

as a function of the reciprocal of absolute temperature.

Figure 11.30 Predominance diagram for the speciation of the chromium(III) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.31 Stability constant (log*

β

1

°) for MnOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.32 Stability constant (log*

β

2

°) for Mn(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.33 Stability constant (log*

β

3

°) for Mn(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 11.34 Stability constant (log*

β

4

°) for Mn(OH)

4

2−

as a function of the reciprocal of absolute temperature.

Figure 11.35 Dependence of log*

β

1

of MnOH

+

on ionic strength in potassium nitrate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.36 Predominance diagram for the speciation of the manganese(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.37 Stability constant (log*

β

1

) for MnOH

2+

as a function of the reciprocal of absolute temperature.

Figure 11.38 Dependence of log*

β

1

of MnOH

2+

on ionic strength in perchloric acid media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.39 Solubility constant (log*

K

s

°) for Fe

3

O

4

(s) as a function of the reciprocal of absolute temperature.

Figure 11.40 Stability constant (log*

β

1

°) for FeOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.41 Stability constant (log*

β

2

°) for Fe(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.42 Stability constant (log*

β

3

°) for Fe(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 11.43 Predominance diagram for the speciation of the iron(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.44 Solubility constant (log*

K

s10

°) for

α

-FeOOH(s) as a function of the reciprocal of absolute temperature.

Figure 11.45 Solubility constant (log*

K

s10

°) for

α

-Fe

2

O

3

(s) as a function of the reciprocal of absolute temperature.

Figure 11.46 Stability constant (log*

β

1

°) for FeOH

2+

as a function of the reciprocal of absolute temperature.

Figure 11.47 Stability constant (log*

β

3

°) for Fe(OH)

3

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.48 Stability constant (log*

β

4

°) for Fe(OH)

4

−

as a function of the reciprocal of absolute temperature.

Figure 11.49 Stability constant (log*

β

22

) for Fe

2

(OH)

2

4+

(1.0 mol l

−1

NaClO

4

) as a function of the reciprocal of absolute temperature.

Figure 11.50 Dependence of log*

β

1

of FeOH

2+

on ionic strength in sodium perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.51 Dependence of log*

β

1

of FeOH

2+

Figure 11.52 Comparison of the calculated stability constants for FeOH

2+

using Eqs. ((11.25)–(11.28)) with those measured for conditions of 4–200 °C and in 0.0025–8.45 mol kg

−1

sodium perchlorate.

Figure 11.53 Dependence of log*

β

2

of Fe(OH)

2

+

on ionic strength in sodium perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.54 Dependence of log*

β

2

of Fe(OH)

2

+

Figure 11.55 Dependence of log*

β

22

of Fe

2

(OH)

2

4+

Figure 11.56 Predominance diagram for the speciation of the iron(III) ion at 25 °C and in 0.01 mol l

−1

NaClO

4

. The solid phase used in the calculations is ferrihydrite (Fe(OH)

3

(s)). Also shown (solid squares) is the measured solubility found under these conditions (Stefansson, 2007). Predominance lines are also shown (dotted lines) assuming a stability constant for Fe(OH)

3

(aq) of log*

β

3

= −13.1 (derived for the conditions stated).

Figure 11.57 Solubility constant (log*

K

s10

°) for β-Co(OH)

2

(s) as a function of the reciprocal of absolute temperature.

Figure 11.58 Stability constant (log*

β

1

°) for CoOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.59 Stability constant (log*

β

2

°) for Co(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.60 Dependence of log*

β

1

of CoOH

+

on ionic strength in sodium perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.61 Predominance diagram for the speciation of the cobalt(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.62 Stability constant (log*

β

1

) for CoOH

2+

(1.0 mol l

−1

NaClO

4

) as a function of the reciprocal of absolute temperature.

Figure 11.63 Dependence of log*

β

1

of CoOH

2+

Figure 11.64 Solubility constant (log*

K

s10

°) for bunsenite (NiO(s)) as a function of the reciprocal of absolute temperature.

Figure 11.65 Solubility constant (log*

K

s10

°) for crystalline theophrastite (

β

-Ni(OH)

2

(s)) as a function of the reciprocal of absolute temperature.

Figure 11.66 Solubility constant (log*

K

s10

°) for microcrystalline theophrastite (

β

-Ni(OH)

2

(s)) as a function of the reciprocal of absolute temperature.

Figure 11.67 Stability constant (log*

β

1

°) for NiOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.68 Stability constant (log*

β

2

°) for Ni(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.69 Dependence of log

β

1

of NiOH

+

Figure 11.70 Dependence of log

β

1

of NiOH

+

Figure 11.71 Dependence of log

β

44

of Ni

4

(OH)

4

4+

Figure 11.72 Predominance diagram for the speciation of the nickel(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.73 Solubility constant (log*

K

s10

°) for cuprite (Cu

2

O(s)) as a function of the reciprocal of absolute temperature.

Figure 11.74 Stability constant (log*

β

1

°) for CuOH(aq) as a function of the reciprocal of absolute temperature.

Figure 11.75 Stability constant (log*

β

2

°) for Cu(OH)

2

−

as a function of the reciprocal of absolute temperature.

Figure 11.76 Solubility constant (log*

K

s10

°) for tenorite (CuO(s)) as a function of the reciprocal of absolute temperature.

Figure 11.77 Stability constant (log*

β

1

°) for CuOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.78 Stability constant (log*

β

2

°) for Cu(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.79 Stability constant (log*

β

4

°) for Cu(OH)

4

2−

as a function of the reciprocal of absolute temperature.

Figure 11.80 Stability constant (log*

β

22

°) for Cu

2

(OH)

2

2+

on the reciprocal of absolute temperature. The uncorrected data from Perrin (1960) are also shown (open circles) with the fit to these data (dashed line). This fit is within the uncertainty of the fit of the remaining data and that of the combined data (solid squares and line).

Figure 11.81 Dependence of log*

β

1

of CuOH

+

Figure 11.82 Dependence of log*

β

2

of Cu(OH)

2

(aq) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.83 Dependence of log*

β

22

of Cu

2

(OH)

2

2+

Figure 11.84 Dependence of log*

β

22

of Cu

2

(OH)

2

2+

on ionic strength in nitrate media and at 20 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.85 Dependence of log*

β

21

of Cu

2

OH

3+

Figure 11.86 Predominance diagram for the speciation of the copper(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 11.87 Solubility constant (log*

K

s10

°) for zincite (ZnO(s)) as a function of the reciprocal of absolute temperature.

Figure 11.88 Solubility constant (log*

K

s10

°) for ε-Zn(OH)

2

(s) as a function of the reciprocal of absolute temperature.

Figure 11.89 Stability constant (log*

β

1

°) for ZnOH

+

as a function of the reciprocal of absolute temperature.

Figure 11.90 Stability constant (log*

β

2

°) for Zn(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 11.91 Stability constant (log*

β

3

°) for Zn(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 11.92 Stability constant (log*

β

4

°) for Zn(OH)

4

2−

as a function of the reciprocal of absolute temperature.

Figure 11.93 Dependence of log*

β

1

of ZnOH

+

Figure 11.94 Dependence of log*

β

1

of ZnOH

+

on ionic strength in NaClO

4

media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.95 Dependence of log*

β

3

of Zn(OH)

3

−

on ionic strength in NaClO

4

media. The solid line is obtained using the derived interaction coefficient and stability constant at zero ionic strength.

Figure 11.96 Dependence of log*

β

4

of Zn(OH)

4

2−

on ionic strength in NaClO

4

(solid squares) and KCl (open circles) media. The solid (NaClO

4

) and dashed (KCl) lines are obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.97 Dependence of log*

β

21

of Zn

2

OH

3+

on ionic strength in chloride (solid squares) and perchlorate (open circles) media. The solid (chloride) and dashed (perchlorate) lines are obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 11.98 Predominance diagram for the speciation of the zinc(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

12 Second and Third Series Transition Metals

Figure 12.1 Solubility data of Meyer

et al.

(1991) and Eriksen

et al.

(1992) for TcO

2

·1.6H

2

O and fitted solubility (solid line) using the solubility constants selected in this review. The open circles are the solubility data of Meyer

et al.

, and the solid squares are from Eriksen

et al.

Figure 12.2 The stability constant (log *

β

1

°) of the formation of RhOH

2+

as a function of the reciprocal of absolute temperature.

Figure 12.3 Solubility constant (log *

K

s10

o

) of Ag

2

O(am) as a function of the reciprocal of absolute temperature.

Figure 12.4 Dependence of log

K

s10

Figure 12.5 Predominance diagram for the speciation of the silver(I) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 12.6 Dependence of log

K

s10

of Cd(OH)

2

(s) (reaction (12.17)) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.7 Dependence of log *

β

1

of CdOH

+

(reaction (2.5), M = Cd

2+

,

p

= 1,

q

= 1) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.8 Dependence of log

β

2

of Cd(OH)

2

(aq) (reaction (2.7), M = Cd

2+

,

p

= 1,

q

= 2) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.9 Dependence of log

β

2

of Cd(OH)

4

2−

(reaction (2.7), M = Cd

2+

,

p

= 1,

q

= 4) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.10 Dependence of log *

β

21

of Cd

2

OH

3+

(reaction (2.5), M = Cd

2+

,

p

= 2,

q

= 1) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.11 Dependence of Δ

H

r

of yellow HgO(s) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 12.12 Dependence of log *

β

1

of HgOH

+

Figure 12.13 Dependence of log *

β

2

of Hg(OH)

2

(aq) on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

13 Aluminium, Gallium, Indium and Thallium

Figure 13.1 Solubility constant (log *

K

s10

°) of gibbsite (Al(OH)

3

(s)) as a function of the reciprocal of absolute temperature.

Figure 13.2 Solubility constant (log *

K

s10

°) of boehmite (γ-AlOOH(s)) as a function of the reciprocal of absolute temperature.

Figure 13.3 Stability constant (log *

β

1

°) of AlOH

2+

as a function of the reciprocal of absolute temperature.

Figure 13.4 Stability constant (log *

β

2

°) of Al(OH)

2

+

as a function of the reciprocal of absolute temperature.

Figure 13.5 Stability constant (log *

β

3

°) of Al(OH)

3

(aq) as a function of the reciprocal of absolute temperature.

Figure 13.6 Stability constant (log *

β

4

°) of Al(OH)

4

−

as a function of the reciprocal of absolute temperature.

Figure 13.7 Stability constant (log *

β

22

) of Al

2

(OH)

2

4+

as a function of the reciprocal of absolute temperature.

Figure 13.8 Stability constant (log *

β

34

) of Al

3

(OH)

4

5+

as a function of the reciprocal of absolute temperature.

Figure 13.9 Stability constant (log *

β

13,32

) of Al

13

(OH)

32

7+

as a function of the reciprocal of absolute temperature.

Figure 13.10 Dependence of log *

K

s10

of gibbsite, Al(OH)

3

(s), on ionic strength in NaCl media at 25 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.11 Dependence of log *

K

s10

of gibbsite, Al(OH)

3

(s), on ionic strength in NaCl media at 100 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.12 Dependence of log *

K

s10

of boehmite, γ-AlOOH(s), on ionic strength in NaCl media at 100 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.13 Dependence of log *

β

22

of Al

2

(OH)

2

4+

Figure 13.14 Dependence of log *

β

34

of Al

3

(OH)

4

5+

Figure 13.15 Dependence of log *

β

13,32

of Al

13

(OH)

32

7+

Figure 13.16 Predominance diagram for the speciation of the aluminium(III) ion at 25 °C and zero ionic strength.

Figure 13.17 Predominance diagram for the speciation of the aluminium(III) ion at 25 °C and 1.0 mol kg

−1

chloride using Al(OH)

3

(am) as the solid species (log *

K

s10

= 10.80).

Figure 13.18 Solubility constant (log *

K

s10

°) of GaOOH(s)) as a function of the reciprocal of absolute temperature.

Figure 13.19 Stability constant (log *

β

1

°) of GaOH

2+

as a function of the reciprocal of absolute temperature.

Figure 13.20 Stability constant (log *

β

2

°) of Ga(OH)

2

+

as a function of the reciprocal of absolute temperature.

Figure 13.21 Stability constant (log *

β

3

°) of Ga(OH)

3

(aq) as a function of the reciprocal of absolute temperature.

Figure 13.22 Stability constant (log *

β

4

°) of Ga(OH)

4

−

as a function of the reciprocal of absolute temperature.

Figure 13.23 Dependence of log *

β

1

of GaOH

2+

on ionic strength in perchlorate media at 25 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.24 Dependence of log *

β

1

of GaOH

2+

on ionic strength in nitrate media at 25 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.25 Dependence of log *

β

2

of Ga(OH)

2

+

on ionic strength in perchlorate media at 25 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.26 Dependence of log *

β

2

of Ga(OH)

2

+

on ionic strength in nitrate media at 25 °C. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.27 Predominance diagram for the speciation of the gallium(III) ion at 25 °C and zero ionic strength.

Figure 13.28 Dependence of log

β

1

of TlOH(aq) on ionic strength in NaClO

4

media. The solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 13.29 Dependence of log

β

2

of Tl(OH)

2

−

on ionic strength in NaClO

4

Figure 13.30 Dependence of log

β

1

of TlOH

2+

on ionic strength in perchlorate media. The solid line is obtained using the derived interaction coefficient and stability constant at zero ionic strength.

14 Tin and Lead

Figure 14.1 Predominance diagram for the speciation of the tin(II) ion at 25 °C. The behaviour in the region of 2 > −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

Figure 14.2 The solubility constant (log *

K

s10

°) of PbO(s) (red) as a function of the reciprocal of absolute temperature.

Figure 14.3 The stability constant (log *

β

1

°) of the formation of PbOH

+

as a function of the reciprocal of absolute temperature.

Figure 14.4 The stability constant (log *

β

2

°) of the formation of Pb(OH)

2

(aq) as a function of the reciprocal of absolute temperature.

Figure 14.5 The stability constant (log *

β

3

°) of the formation of Pb(OH)

3

−

as a function of the reciprocal of absolute temperature.

Figure 14.6 Dependence of log

β

1

of PbOH

+

on ionic strength in NaClO

4

Figure 14.7 Dependence of log

β

1

of PbOH

+

on ionic strength in nitrate media (solid squares for 25 °C and open circles for 18 °C). The solid (25 °C) and dashed (18 °C) lines are obtained using the derived interaction coefficients and stability constant at zero ionic strength.

Figure 14.8 Dependence of log *

β

2

of Pb(OH)

2

(aq) on ionic strength in NaClO

4

Figure 14.9 Dependence of log

β

3

of Pb(OH)

3

−

on ionic strength in NaClO

4

Figure 14.10 Dependence of log *

β

21

of Pb

2

OH

+

on ionic strength in NaClO

4

(squares) and nitrate (circles: 25 °C and triangles: 18 °C) media. The solid line (perchlorate) is obtained using the derived interaction coefficient and stability constant at zero ionic strength. The dashed line (nitrate: 25 °C) is obtained using the stability constant obtained at zero ionic strength from the perchlorate data and the derived interaction coefficient from the nitrate data. The dotted line (nitrate: 18 °C) is obtained from the derived interaction coefficient and stability constant at zero ionic strength.

Figure 14.11 Dependence of log

β

34

of Pb

3

(OH)

4

2+

on ionic strength in NaClO

4

and nitrate media. The dashed (nitrate) line has been obtained using the derived interaction coefficient and stability constant at zero ionic strength. The solid (NaClO

4

) line has been obtained from the zero ionic strength stability constant calculated in nitrate media and the derived ion interaction coefficients for NaClO

4

media.

Figure 14.12 Dependence of log *

β

44

of Pb

4

(OH)

4

4+

on ionic strength in NaClO

4

(squares) and nitrate (circles: 25 °C and triangles: 18 °C) media. The solid line (perchlorate) is obtained using the derived interaction coefficient and stability constant at zero ionic strength. The dashed line (nitrate: 25 °C) is obtained using the stability constant obtained at zero ionic strength from the perchlorate data and the derived interaction coefficient from the nitrate data. The dotted line (nitrate: 18 °C) is obtained from the derived interaction coefficient and stability constant at zero ionic strength.

Figure 14.13 Dependence of log

β

68

of Pb

6

(OH)

8

4+

on ionic strength in NaClO

4

and nitrate media. The solid (NaClO

4

) and dashed (nitrate) lines have been obtained from the zero ionic strength stability constant and ion interaction coefficients calculated in the separate media.

Figure 14.14 Predominance diagram for the speciation of the lead(II) ion at 25 °C. The behaviour in the region of −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

15 Bismuth and Polonium

Figure 15.1 Solubility constant (log *

K

s10

°) as a function of the reciprocal of absolute temperature.

Figure 15.2 Stability constant (log *

β

2

°) for Bi(OH)

2

+

as a function of the reciprocal of absolute temperature.

Figure 15.3 Stability constant (log *

β

3

°) for Bi(OH)

3

(aq) as a function of the reciprocal of absolute temperature.

Figure 15.4 Dependence of log *

β

1

of BiOH

2+

on ionic strength in perchlorate media (solid squares – the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). Also shown are the data from Antonovich

et al.

(1975) obtained in nitrate media (open circles – the dashed line is also obtained using the extended specific ion interaction theory).

Figure 15.5 Dependence of log *

β

2

of Bi(OH)

2

+

on ionic strength in perchlorate media (solid squares – the solid line is obtained using the derived interaction coefficients and stability constant at zero ionic strength). Also shown are the data from Antonovich

et al.

(1975) obtained in nitrate media (open circles – the dashed line is also obtained using the extended specific ion interaction theory).

Figure 15.6 Dependence of log *

β

4

of Bi(OH)

4

−

Figure 15.7 Dependence of log *

β

6,12

of Bi

6

(OH)

12

6+

Figure 15.8 Predominance diagram for the speciation of the bismuth(III) ion at 25 °C. The behaviour in the region of 2 > −log [H

+

] > 12 should be treated with caution due to changes in activity coefficients.

16 Prediction of Stability and Solubility Constants

Figure 16.1 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

1

°) of MOH

(

z

−1)

species.

Figure 16.2 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

2

°) of M(OH)

2

(

z

−2)

species.

Figure 16.3 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

3

°) of M(OH)

3

(

z

−3)

species.

Figure 16.4 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

4

°) of M(OH)

4

(

z

−4)

species.

Figure 16.5 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

22

°) of M

2

(OH)

2

(2

z

−2)

species.

Figure 16.6 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

34

°) of M

3

(OH)

4

(3

z

−4)

species.

Figure 16.7 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

35

°) of M

3

(OH)

5

(3

z

−5)

species.

Figure 16.8 Comparison of predicted and measured stability constants (log *

β

pq

°) of M

p

(OH)

q

(

pz

−

q

)

species.

Figure 16.9 Relationship between the function

z

[14.52 − 0.1408

g

1

(

z

/

r

2

+

g

2

)] and the solubility constants (log *

K

s10

°) of metal oxide/hydroxidephases.

Figure 16.10 Relationship between the stability and solubility constants for M(OH)

6

2−

species and MO

2

(s) phases.

Figure 16.11 Relationship between the function

g

1

(

z

/

r

2

+

g

2

) and the stability constants (log *

β

1

°) of MOH

(

z

−1)

species at 100 °C.

Volume 1

2 Theory

Table 2.1 Calculated values of

A

and

B

in Eqs. (2.44) and (2.45)

Table 2.2 Calculated values of the Debye–Hückel parameters

A

and

B

using a multiple regression equation

Table 2.3 Comparison of calculated and literature values of the ion size parameter at the saturation pressure of water