Thermodynamics of Reactions

The following sections are included:

• Introduction

• Standard Free Energies of Formation

• Predominance Area Diagram for the Fe–H₂O System, 298.15 K

  • Reactions Involving Dissolved Species and Solids, Nonredox

  • Reactions Involving Oxidation and Reduction

  • The Hydrogen Scale

  • The pe Scale

  • Equilibria between Solids

  • Equilibria between Ions and Solids

  • Equilibria between Ions

  • Representation of Equilibria

  • Discussion of the Diagram, Fig. 3.1

• The Sulphur–Water System, 298.15 K

  • Data Used

  • Acid Dissociation Constants

  • Redox Equilibria

  • Discussion of the Diagram, Fig. 3.2

• The Copper–Sulphur–Water System, 298.15 K

  • Data Used

  • Equilibria

  • Discussion of the Diagram, Fig. 3.3

• Thermodynamics of Electrolyte Solutions at Elevated Temperatures

  • Construction of Predominance Area Diagrams for Elevated Temperatures

  • The S–H₂O System, 423 K

  • The Cu–S–H₂O System, 423 K

  • The Fe–S–H₂O System, 423 K

  • Estimation of Entropies and Heat Capacities of Ionic Species

• Precipitation of Metals from Solution by Reduction

  • Introduction

  • Deposition of Metals by Electrolysis

  • Precipitation of Metals by Hydrogen Gas

• Precipitation of Compounds from Solution

  • Introduction

  • The Second Dissociation Constant of H₂S(aq)

  • Precipitation of Sparingly Soluble Metal Sulphides

  • Nucleation and Crystal Growth

  • Precipitation of Arsenates

  • Precipitation of Iron as Oxides

• pH at Elevated Temperatures

• Precipitation of Solids at Elevated Temperatures

  • Introduction

  • True Solubility and Temperature

  • Hydrolysis and Temperature

• Speciation Diagrams

  • Introduction

  • Speciation in Multicomponent Solutions

  • Precipitation of Trace Metals in Estuarine Water

  • Adsorption and Coprecipitation of Metals from Solutions