Pyrite Oxidation and Its Control

V. P. Evangelou

• Publisher: CRC Press
• Year: 2018
• Print ISBN: 9780849347320
• eBook ISBN: 9781351420792
• Topic: Physical Sciences
• Subtopic: Chemistry
• Index: Chemistry

Pyrite Oxidation and its Control is the single available text on the market that presents the latest findings on pyrite oxidation and acid mine drainage (AMD). This new information is an indispensable reference for generating new concepts and technologies for controlling pyrite oxidation. This book focuses on pyrite oxidation theory, experimental findings on oxidation mechanisms, as well as applications and limitations of amelioration technologies. The text also includes discussions on the theory and potential application of novel pyrite microencapsulation technologies for controlling pyrite oxidation currently under investigation in the author's laboratory.

PART I
REVIEW OF BASIC CONCEPTS OF SURFACE CHEMISTRY AND CHEMISTRY AND CONTROL OF ACID DRAINAGE

Chapter 1

A REVIEW ON SOLUTION/MINERAL CHEMISTRY

1.1 Ion Activity versus Concentration

It is generally agreed that chemical reactions in nature (e.g., mineral dissolution/precipitation) take place because of electrical and chemical gradients. With regard to the former, there is disagreement as to how one may measure these gradients since they are not directly related to concentration (Amacher, 1984; Sposito, 1984; Sposito and Mattigod, 1979). Aided by computer technology, approaches have been introduced that allow estimation of single-ion activity (a direct reflection of the chemical potential of ionic species) in mixed electrolyte systems (Davies, 1962; Adams, 1971; Sposito, 1984). Experimental data (Pavan and Bingham, 1982; Sposito, 1984) support the argument that single-ion activity best describes/predicts chemical reactions.

Single-ion activity in solution is estimated by employing the equation of effective ionic strength (I)

$I=1/2{\displaystyle \sum _{\text{j}=1}^{\text{n}}{Z}_{\text{j}}^2{m}_{\text{j}}}$

(1.1)

where Z_j and m_j are the charge and effective molar concentration of ionic species (Sposito, 1984). Based on I, activity coefficients of ionic species can be calculated using the Davies equation (Davies, 1962):

$\log {\text{γ}}_{\text{j}}=-A{Z}_{\text{j}}^2\left[I^{1/2}/\left(1+I^{1/2}\right)-0.3I\right]$

(1.2)

where γ is the activity coefficient of ionic species j, and A is a constant equal to 0.512. Single ion activity (α_j) can then be estimated by:

${\text{α}}_{\text{j}}={\text{γ}}_{\text{j}}{\text{m}}_{\text{j}}$

(1.3)

The above calculations necessitate the use of ion association models to make estimates of concentration of ionic species (m_j) using as inputs elemental concentrations (see section 1.1.3, Ion pairing) (Adams, 1971; Sposito and Mattigod, 1979; Sposito, 1984; Evangelou, 1986). However, direct measurement of single-ion activity in mixed electrolyte systems for the purpose of verifying ion association models cannot be made (Amacher, 1984; Sposito, 1984); validity of ion-association models can only be demonstrated indirectly.

In order to gain an appreciation of the importance of single-ion activity in describing/predicting chemical reactions as opposed to elemental concentration, one needs to look into the effect of ionic strength and ion valence on single-ion activity coefficients (γ) (Equation 1.2). This is demonstrated in Figure 1.1. It shows that at an ionic strength of 0.04 M, the γ value for say Ca²⁺, is approximately 0.5, while that of a monovalent ion is approximately 0.8.

Figure 1.1. Relationship between solution ionic strength and activity coefficients of ions of different valencies. Calculated utilizing the Davies equation (Davies, 1962). (From Evangelou and Sobek, 1988. With permission.)

Examples of practical implications of single-ion activity with respect to pH and mineral solubility are given in sections 1.1.1 and 1.1.2.

1.1.1 Solution pH

Most reactions in soil/geologic systems are pH driven because of the hydrogen’s (H⁺) great potential to react with solution chemical species and/or mineral surfaces due to its small size and high electron accepting potential.

Technically, pH is the negative logarithm of the hydrogen ion activity (Equations 1.4)

$\text{pH}=-\log \left({\text{H}}^{+}\right)$

(1.4)

The pH scale ranges from 0 to 14, with 7 being the neutral point. Values greater than 7 are basic or alkaline, while values less than 7 are acidic. For example, if the hydrogen ion activity of a water sample is 10⁻⁷ mol L⁻¹, its pH is 7, since 7 is the negative logarithm of 10⁻⁷.

Water emanating from surface mine sites for example can vary in ionic strength (I) anywhere from 0.001 to 0.1. Values greater than 0.1 can be observed in acid mine drainages (Lekhakul, 1981). For a solution with an I of 0.008 M, the γ value for H⁺ will be 0.9. On the other hand, a solution with an I value of 0.04 M will exhibit a γ value for H⁺ of 0.8. Since pH represents H⁺ activity, the same pH in the above systems would represent a different H⁺ concentration. For example, at pH of 3 and at an I value of 0.008 M:

$\left[{\text{H}}^{+}\right]=1.11\times {10}^{-3}{\text{mol L}}^{-1}$

(1.5)

while at an I of 0.04 M

$\left[{\text{H}}^{+}\right]=1.25\times {10}^{-3}{\text{mol L}}^{-1}$

(1.6)

Therefore, at an ionic strength of 0.008 M the H⁺ activity represents 90% of its concentration, while at an ionic strength of 0.04 M, the H⁺ activity represents 80% of its concentration.

The above example demonstrates that even in the case of a monovalent ion...

Table of contents

• Cover
• Title Page
• Copyright Page
• Table of Contents
• INTRODUCTION
• PART I REVIEW OF BASIC CONCEPTS OF SURFACE CHEMISTRY AND CHEMISTRY AND CONTROL OF ACID DRAINAGE
• PART II MOLECULAR CHEMISTRY OF PYRITE OXIDATION KINETICS AND PYRITE REDOX CHEMISTRY
• PART III ROLE OF BACTERIA ON PYRITE OXIDATION
• PART IV ADVANCES ON PYRITE MICROENCAPSULATION TECHNOLOGIES
• PART V SUMMARY
• REFERENCES
• SUBJECT INDEX