A Problem-Solving Approach to Aquatic Chemistry E-Book

James N. Jensen

• Herausgeber: John Wiley & Sons
• Kategorie: Fachliteratur
• Sprache: Englisch

Seitenzahl: 1405

Veröffentlichungsjahr: 2022

Beschreibung

A Problem-Solving Approach to Aquatic Chemistry Enables civil and environmental engineers to understand the theory and application of aquatic equilibrium chemistry The second edition of A Problem-Solving Approach to Aquatic Chemistry provides a detailed introduction to aquatic equilibrium chemistry, calculation methods for systems at equilibrium, applications of aquatic chemistry, and chemical kinetics. The text directly addresses two required ABET program outcomes in environmental engineering: "... chemistry (including stoichiometry, equilibrium, and kinetics)" and "material and energy balances, fate and transport of substances in and between air, water, and soil phases." The book is very student-centered, with each chapter beginning with an introduction and ending with a summary that reviews the chapter's main points. To aid in reader comprehension, important terms are defined in context and key ideas are summarized. Many thought-provoking discussion questions, worked examples, and end of chapter problems are also included. Each part of the text begins with a case study, a portion of which is addressed in each subsequent chapter, illustrating the principles of that chapter. In addition, each chapter has an Historical Note exploring connections with the people and cultures connected to topics in the text. A Problem-Solving Approach to Aquatic Chemistry includes: * Fundamental concepts, such as concentration units, thermodynamic basis of equilibrium, and manipulating equilibria * Solutions of chemical equilibrium problems, including setting up the problems and algebraic, graphical, and computer solution techniques * Acid-base equilibria, including the concepts of acids and bases, titrations, and alkalinity and acidity * Complexation, including metals, ligands, equilibrium calculations with complexes, and applications of complexation chemistry * Oxidation-reduction equilibria, including equilibrium calculations, graphical approaches, and applications * Gas-liquid and solid-liquid equilibrium, with expanded coverage of the effects of global climate change * Other topics, including chemical kinetics of aquatic systems, surface chemistry, and integrative case studies For advanced/senior undergraduates and first-year graduate students in environmental engineering courses, A Problem-Solving Approach to Aquatic Chemistry serves as an invaluable learning resource on the topic, with a variety of helpful learning elements included throughout to ensure information retention and the ability to apply covered concepts in practical settings.

Leseprobe

A Problem‐Solving Approach to Aquatic Chemistry

Second Edition

James N. JensenUniversity at Buffalo, NY, US

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Library of Congress Cataloging‐in‐Publication DataNames: Jensen, James N., author.Title: A problem‐solving approach to aquatic chemistry / James N. Jensen, SUNY—Buffalo, NY, US.Description: Second edition. | Hoboken : Wiley, [2023] | Includes bibliographical references and index.Identifiers: LCCN 2022022564 (print) | LCCN 2022022565 (ebook) | ISBN 9781119884347 (cloth) | ISBN 9781119884354 (adobe pdf) | ISBN 9781119884361 (epub)Subjects: LCSH: Water chemistry. | Environmental chemistry.Classification: LCC GB855 .J46 2023 (print) | LCC GB855 (ebook) | DDC 551.4801/54—dc23/eng20221004LC record available at https://lccn.loc.gov/2022022564LC ebook record available at https://lccn.loc.gov/2022022565

Cover Design: WileyCover Image: © Nick Brundle Photography/Getty Images

PREFACE

FOR THE STUDENT

It is my honor and privilege to help you on your journey through aquatic chemistry with the second edition of this book. I hope that this text will deepen your appreciation of water. As with the first edition, my goal is that you come to look at rain, a stream, or tap water and admire the wealth of chemistry in every drop.

Each chapter of the text is built to help you learn. Overall, the text is designed to mimic the Socratic method of teaching. You will encounter Thoughtful Pause1 boxes throughout the text. Please stop and try to answer the questions in the Thoughtful Pauses before continuing to read. You will find over 215 important definitions and over 300 Key Ideas highlighted in the margins and summarized at the end of each chapter. Do you learn better through examples? The text contains over 180 worked examples.

To help you integrate the material, each part of the text contains a case study that is developed through several chapters. The last chapter of the text contains five integrative case studies. These are large, practical problems that combine many of the tools you will learn to use throughout the text.

The first edition of this text was accompanied by a user‐friendly equilibrium calculation app called Nanoql. In this edition, the app has been ported to Excel as Nanoql SE (spreadsheet edition). The text provides instructions on how to use Nanoql SE, along with a number of examples.

To give you an appreciation for the fascinating history of the ideas presented in the text, the number of Historical Notes in the second edition has been greatly expanded. You will travel from the battlefields of World War II to deep‐sea trenches. You will meet famous long‐bearded chemists and an underappreciated female scientist who provided the first experimental evidence for global climate change. Along the way, you will explore the connections between aquatic chemistry and Sherlock Holmes, Lewis Carroll, and Kurt Vonnegut.

And, of course, what textbook on water chemistry would be complete without haikus? Each part of the text is summarized in seventeen syllables to aid in reinforcing the main ideas.

FOR THE INSTRUCTOR

The text continues to be appropriate for a one‐semester course in water chemistry (or in an environmental chemistry course emphasizing aquatic systems) for advanced undergraduates, first‐year graduate students, or self‐studying professionals. This edition features extended coverage of the effects of global climate change on the aquatic environment and surface complexation modeling. The number of end‐of‐chapter problems has been nearly doubled to over 340 problems in this edition. In addition, the Nanoql equilibrium calculation app has been updated to a more accessible Excel‐based version. Numerous other spreadsheet solutions are developed in the text.

ACKNOWLEDGMENTS

Once again, I wish to acknowledge my great debt to the aquatic chemistry teachers I have had. Since the first edition, we lost Jim Morgan and Phil Singer. Their legacy lives on in countless labs, treatment plants, and classrooms through the world, and also in the pages of this book. I also am indebted to my colleagues at the University at Buffalo for their steadfast support. They have made me a better teacher. Numerous comments from my students and other readers of the first edition have influenced and improved the text.

I am deeply appreciative to the staff at Wiley for their guidance and support. In particular, I am thankful for the efforts of Paul Sayer, Kimberly Monroe‐Hill, Cheryl Ferguson, Michael New, and Judit Anbu Hena Daniel in bringing this book to fruition.

James N. JensenBuffalo, NYOctober 2022

Note

1

Professor James J. Morgan (who you will meet throughout the text, but especially in

Chapter 24

), referred to them as “Thoughtful Paws” boxes.

PART IFundamental Concepts

Chapter 1:

Getting Started with the Fundamental Concepts

Chapter 2:

Concentration Units

Chapter 3:

Thermodynamic Basis of Equilibrium

Chapter 4:

Manipulating Equilibrium Expressions

The fundamental things apply

As time goes by.

–Herman Hupfeld (

1931)

Part I Haiku

Use molar units: Equilibrium constants are dimensionless.

…the beginning is the most important part of any work…

–Plato (c. 375

BCE

)

CHAPTER 1Getting Started with the Fundamental Concepts

1.1 INTRODUCTION

The first part of this text reviews fundamental concepts that must be mastered prior to learning how to calculate and interpret species concentrations in aquatic systems. In this chapter, the motivation for studying chemical species and a few general principles concerning aquatic systems are presented.

In Section 1.2, the motivation for why engineers and scientists are interested in individual chemical species concentrations at equilibrium will be discussed. Important water quality parameters, called primary variables, are introduced in Section 1.3. It is impossible to study water chemistry without a little knowledge of the structure of water. A few of the unique properties of water will be explored in Section 1.4, especially as they relate to the chemical reactions that occur in water. In Section 1.5, a road map for Part I of the text is presented and discussed. Finally, the Part I case study is presented at the conclusion of this chapter. Before beginning Part I of the text, you are urged to review the chemistry background material in Appendix A (Section A.2).

1.2 WHY CALCULATE CHEMICAL SPECIES CONCENTRATIONS AT EQUILIBRIUM?

1.2.1 Overview

The bulk of this book is dedicated to the calculation of species concentrations at equilibrium. The focus here is on chemical species that undergo chemical reactions; in other words, reactive species. More specifically, the emphasis here is on chemical species which react with water. Reactions with water are called hydrolysis reactions (from the Greek hydōr water + lyein to loosen). When substances react with water, numerous other compounds can be formed. Indeed, the richness of aquatic chemistry stems from the large number of substances that react not only with water but also with the products of myriad other hydrolysis reactions.

hydrolysis reactions:

reactions with water

Key idea: Hydrolysis reactions produce a wealth of dissolved chemical species

This richness is illustrated in Figure 1.1. Inputs of chemical species (from aqueous discharge, runoff, atmospheric deposition, and dissolution from sediments) react with water to form hydrolysis products. The hydrolysis products and input chemicals react further to increase the complexity of aquatic systems.

FIGURE 1.1 Complexity of Aquatic Systems

(rain cloud image: OpenClipart‐Vectors/Pixabay)

So why so much interest in calculating the equilibrium concentrations of chemical species? This question is really two questions. First, why calculate the concentrations of individual chemical species? Second, why calculate species concentrations at equilibrium?

1.2.2 Importance of Individual Chemical Species

Throughout this text, you will see that knowing the concentrations of individual chemical species is critically important in analyzing many environmental problems. At first glance, this statement may not make sense. After all, many environmental regulations are based on total concentrations of classes of compounds rather than on the concentrations of individual species. Should you be more concerned about the total amount of mercury or phenol or ammonia than about individual species stemming from the hydrolysis of mercury, phenol, or ammonia?

Key idea: The ability to calculate the concentrations of individual chemical species is critically important in analyzing many environmental problems

In fact, you will find that individual species frequently are more important. Three general examples will illustrate this point. First, adverse impacts on human health and ecosystem viability may be due to only one or several of a large number of related hydrolysis products. A prime example is the transition metals (such as mercury, copper, zinc, cadmium, iron, and lead), in which toxicity varies dramatically among the hydrolysis products. Another example is cyanide. Hydrogen cyanide (HCN) is much more toxic to humans than cyanide ion (CN–).

Second, the success of engineered treatment systems may depend on knowledge of the concentrations of key individual species. Since hydrolysis products vary in their physical, chemical, and biochemical properties, the design and operation of treatment processes depend on quantitative models for the concentrations of individual chemical species. For example, the addition of gaseous chlorine to wastewater for disinfection results in the formation of many chemical species (including HOCl, OCl–, NH2Cl, and NHCl2), each of which differs in its ability to inactivate (i.e., kill) microorganisms.

Third, individual species vary greatly in how readily they cross cell membranes or cell walls and are assimilated by aquatic biota. Thus, understanding the cycling of trace nutrients in the aquatic environment (and humankind's impact on nutrient cycling) requires knowledge of the concentrations of individual chemical species.

As an example of the importance of the concentrations of individual chemical species, consider the soup created when copper sulfate crystals, CuSO4(s),1 are added to a reservoir for algae control. The CuSO4(s) dissolves in water to form a copper‐containing ion (called the aquo cupric ion) and sulfate. The structure of the aquo cupric ion is usually abbreviated as Cu2+. The Cu2+ ions thus formed react very quickly with water to form a number of hydrolysis products, including CuOH+, Cu(OH)2(aq), Cu(OH)3–, Cu(OH)42–, and Cu2(OH)22+. Under certain chemical conditions, copper may precipitate as CuO(s). As you spread the copper sulfate from the back of a boat, carbon dioxide in the atmosphere is equilibrating with the reservoir water to form its own hydrolysis products. The hydrolysis products of carbon dioxide are H2CO3, HCO3–, and CO32–. The aquo cupric ion will react to some extent with the hydrolysis products of carbon dioxide to form CuCO3(aq), Cu(CO3)22–, and perhaps even solids containing copper and carbonate (CO32–). By adding one copper compound to a natural water body, you may be faced with accounting for as many as 10 copper‐containing species even in a relatively simple chemical model.

Of course, the real world is even more complex. The reservoir water contains many more species that can react with copper than just hydroxide (OH–) and carbonate. A realistic model for copper in the reservoir would have to include the reactions of Cu2+ with (among other chemical species) chloride, amino acids, ammonia, particulates, and microorganisms. In reality, the act of throwing copper sulfate crystals into the reservoir will produce dozens of chemical species containing copper.

Key idea: Doses depend on both the required concentration of the target individual species and the chemistry of the water

Why should you care that copper sulfate forms many copper‐containing species in a lake? Remember that copper is added to kill algae. It is well‐established that copper toxicity to algae is due almost entirely to one chemical species: Cu2+ (Jackson and Morgan 1978). Thus, to determine the copper sulfate dose, you must be able to calculate the concentration of Cu2+ after a certain amount of copper sulfate is added to the reservoir. Since the Cu2+ concentration usually is exceedingly small, this is akin to counting needles of Cu2+ in this haystack of copper‐containing species. In practice, you would back‐calculate the copper sulfate dose required to achieve a required level of Cu2+. Even if two reservoirs had the same amounts of algae, different water chemistries in the reservoirs may lead to very different copper sulfate doses to achieve the same Cu2+ concentration. The chemistry of the water determines how the required concentration of one species (Cu2+) is translated back into a copper sulfate dose.

The process of relating a dose to a required concentration of an individual chemical species is illustrated in Figure 1.2. The arrows in Figure 1.2 indicate chemical reactions that must be included in a mathematical model to allow for the determination of the copper sulfate dose. In this text, you will learn the tools to make quantitative decisions to solve similar problems in the aquatic environment.

FIGURE 1.2 Qualitative Relationship Between the Dose Required and End Species Concentrations Desired

1.2.3 Importance of Equilibrium

An entire chapter of this book (Chapter 3) is devoted to developing the thermodynamic basis of equilibrium. For the present, you can think of the equilibrium state as the condition in which the concentrations of all chemical species do not change with time. To impose equilibrium on a chemical system, the interesting and important time‐dependent nature of chemical concentrations are excluded. The study of the rates of chemical reaction is called chemical kinetics and is covered in Chapter 23. Why constrain the discussion mainly to the equilibrium state here, with shorter coverage of chemical kinetics?

chemical kinetics: the study of chemical reaction rates

There are two reasons for focusing on equilibrium. First, many of the chemical reactions you will examine in this text are fast. For example, the reaction of H+ and OH– to form water occurs on the time scale of 10–5 s at natural water conditions (Morgan and Stone 1985). Second, equilibrium models give insight into chemical systems, even when kinetics are known to be important. As an example, equilibrium chemistry provides a good framework to understand coagulation chemistry in drinking water treatment. Chemical kinetics also play a role in coagulation, but equilibrium models remain useful. Thus, the focus on equilibrium is somewhat constraining but not overly restrictive.

Key idea: The equilibrium state is a useful approximation of many aquatic chemical systems

1.3 PRIMARY VARIABLES: IMPORTANCE OF pH AND pe

As discussed in Section 1.2, it is important to know the concentrations of individual chemical species in aquatic systems. Some species may be of greater or lesser importance, but a small number of chemical species often control the chemistry of an aquatic system. “Controlling the chemistry” means that certain chemical species play a dominant role in determining the concentrations of other chemical species.

The concentrations of the controlling species are expressed in parameters called primary variables.2 The two most important primary variables in water are pH and pe. The primary variable pH is defined as the negative of the log of the activity of the ion H+.3 Similarly, pe is defined as the negative of the log of the activity of the electron, e–. (See the Historical Note at the end of this chapter for the story of what the p in pH and pe means.) Since the hydrogen atom contains one proton and one electron in its most abundant state, H+ is a proton. Thus, the transfer of H+ is sometimes called proton transfer or acid–base chemistry. Electron transfer is also called oxidation‐reduction chemistry or redox chemistry. Acid–base chemistry is discussed in more detail in Chapter 11. Redox chemistry is discussed in more detail in Chapter 16.

primary variable: a species concentration (or function of a species concentration) that controls the chemistry of a system

pH: a primary variable controlling H+ (proton) transfer and defined by –log(H+ activity)

pe: a primary variable controlling electron transfer and defined by –log(e– activity)

As noted in Appendix A, a low pH corresponds to high activity of H+. Conditions of high H+ activity are called acidic conditions. Thus, low pH (high H+ activity) means acidic conditions. High pH (low H+ activity) corresponds to basic conditions. Low pe (high e– activity) corresponds to reducing conditions, and high pe (low e– activity) corresponds to oxidizing conditions.

1.4 PROPERTIES OF WATER

1.4.1 Introduction

Water is a unique substance. In freshman chemistry courses, you probably learned of the periodic nature of the properties of the elements. Elements in the same column of the periodic table share similar chemical properties. Yet water, H2O, has very different physical and chemical properties than other dihydrogen complexes of elements in oxygen's column of the periodic table. For example, water is a liquid under standard temperature and pressure conditions, but H2S is a gas.

The unique properties of water stem from the large differences between the affinity of oxygen and hydrogen for electrons. Oxygen has a much higher affinity for electrons than hydrogen, resulting in a highly polarized bond between O and H. In fact, the oxygen atoms in water are partially negatively charged and the hydrogen atoms are partially positively charged, as illustrated in Figure 1.3. As a result of this polar bond and the difference in partial charge, oxygen atoms in water have the ability to form weak, but important, chemical bonds with more than two hydrogen atoms. These weak bonds are called hydrogen bonds.

hydrogen bonds: weak bonds formed between an electronegative atom (e.g., oxygen) bonded to an H atom and another H atom

FIGURE 1.3 Charge Distribution among the Atoms in Water

Without exaggeration, hydrogen bonds influence every molecular interaction in the aquatic environment. Consider, for example, the influence of hydrogen bonds on ice.4 The structure of ice is different than the structure of water. In ice, all the water molecules participate in four hydrogen bonds arranged in a tetrahedral shape. The shape results in an open structure for ice and a corresponding density of 0.92 g/cm3. The density of ice is less than the density of liquid water at 0°C, so ice floats. Hydrogen bonds reduce the O‐H bond length by only about 4% in ice compared to water. This small change is responsible for the open structure and low density of ice. If not for this impact of hydrogen bonds, ice would sink and water bodies would have frozen solid during the ice ages. If hydrogen bonds were slightly different, life on Earth might not have survived long enough to allow you to read this book.

Hydrogen bonds affect the three‐dimensional shape of water. The bond distance between oxygen in water and the hydrogen‐bonded hydrogen is about twice the bond distance between the oxygen in water and one of its own hydrogen atoms. Hydrogen bonding allows for the formation of very large clusters of water molecules (about 400 molecules per cluster; Luck 1998). The clusters are shifting constantly since the average lifetime of a hydrogen bond is only a few picoseconds (10–12 s). The shifting clusters of water molecules contribute to the solubilization of ions in water, as will be discussed in Section 1.4.2.

1.4.2 Solubility of Ionic Species

Throughout this text, it will be assumed that salts containing sodium, potassium, or chloride at low concentrations dissolve nearly completely in water to form ions. Why are salts so soluble in water? Hydrogen bonding allows for the orientation of a large number of molecules simultaneously. This orientation reduces the field strength of an applied field significantly.

More specifically, the orientation reduces the attractive forces between pairs of anions and cations. As the attractive forces are reduced, the salt can be solubilized. For example, you can show that water significantly reduces the attractive forces between Na+ and Cl– (see Worked Example 1.1).

Worked Example 1.1 Solubilization of Sodium Chloride in Water

Why does NaCl dissolve in water?

Solution

One reason NaCl dissolves in water is that the potential energy of interaction between Na+ and Cl– is reduced in water. To illustrate this point, calculate the potential energy of interaction between Na+ and Cl– separated by 1 nm in water.

The potential energy of two particles of charge q1 and q2 separate by a distance r is given by Coulomb's law:

where 0 is the permittivity of a vacuum (= 8.854×10–12 J–1C2m–1) and is the relative permittivity or dielectric constant (= 1 for a vacuum and 80 for water). The unit charge on Na+ is +1.602×10–19 C. It is

–1.602×10–19 C on Cl–. (The term is Coulomb's constant, ke, after Charles‐Augustin de Coulomb, 1736–1806.)

Plugging in the values, the potential energy is –139 kJ per mole in a vacuum and –1.7 kJ per mole in water. Here, the negative sign indicates an energy of attraction. Recall that one mole is 6.022×1023 ions).

The attractive potential energy of interaction between Na+ and Cl– in water is very small (only about one‐tenth of the energy of a hydrogen bond).

Sodium chloride is soluble in water, mainly because the attractive energy between Na+and Cl–is reduced by water.