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Professor George W. Luther, III, School of Marine Science & Policy, University of Delaware, USA Professor Luther has joint appointments in the Department of Chemistry and Biochemistry, Department of Civil and Environmental Engineering and the Department of Plant and Soil Science. Professor Luther taught an ACS accredited course on advanced inorganic chemistry from 1973-1986 to senior undergraduate students. As he moved into environmental and marine chemistry, he began using environmental examples in inorganic chemistry. In 1988, he started a similar course titled 'Marine Inorganic Chemistry' that has been taught biannually at the University of Delaware, attracting students in Chemical Oceanography, Chemistry and Biochemistry, Geology / Geochemistry, Civil and Environmental Engineering and Plant and Soil Science. In 2013, Professor Luther was awarded the Geochemistry Division Medal by the American Chemical Society for his wide-ranging contributions to aqueous geochemistry. He is recognised for the application of physical inorganic chemistry to the transfer of electrons between chemical compounds in the environment, and also the development of chemical sensors for quantifying the presence of elements and compounds in natural waters. Professor Luther was named a fellow of the American Association for the Advancement of Science in 2011 and the American Geophysical Union in 2012.
About the Author xv
Preface xvii
Companion Website xix
1. Inorganic Chemistry and the Environment 1
1.1 Introduction 1
1.1.1 Energetics of Processes 1
1.2 Neutron-Proton Conversion 3
1.3 Element Burning Reactions - Buildup of Larger Elements 4
1.4 Nuclear Stability and Binding Energy 5
1.4.1 The "r" and "s" Processes 6
1.5 Nuclear Stability (Radioactive Decay) 8
1.6 Atmospheric Synthesis of Elements 8
1.7 Abundance of the Elements 8
1.7.1 The Cosmos and the Earth's Lithosphere 8
1.7.2 Elemental Abundance (Atmosphere, Oceans, and Human Body) 10
1.8 Scope of Inorganic Chemistry in Geochemistry and the Environment 17
1.8.1 Elemental Distribution Based on Photosynthesis and Chemosynthesis 17
1.8.2 Stratified Waters and Sediments - the Degradation of Organic Matter by Alternate Electron Acceptors 19
1.9 Summary 21
1.9.1 Environmental Inorganic Chemistry 22
References 22
2. Oxidation-Reduction Reactions (Redox) 24
2.1 Introduction 24
2.1.1 Energetics of Half Reactions 24
2.1.2 Standard Potential and the Stability of a Chemical Species of an Element 26
2.2 Variation of Standard Potential with pH (the Nernst Equation) 29
2.3 Thermodynamic Calculations and pH Dependence 29
2.4 Stability Field of Aqueous Chemical Species 31
2.5 Natural Environments 32
2.6 Calculations to Predict Favorable Chemical Reactions 32
2.6.1 Coupling Half-Reactions 34
2.6.2 One-Electron Oxygen Transformations with Fe2+ and Mn2+ to Form O2- 35
2.7 Highly Oxidizing Conditions 38
2.7.1 Ozonolysis Reactions 38
2.7.2 Atmospheric Redox Reactions 39
Appendix 2.1 Gibbs Free Energies of Formation 43
References 43
3. Atomic Structure 45
3.1 History 45
3.2 The Bohr Atom 46
3.3 The Schrodinger Wave Equation 47
3.4 Components of the Wave Function 50
3.4.1 Radial Part of the Wave Function, R(r) 50
3.4.2 Angular Part of the Wavefunction Ylml(¿¿¿¿, ¿¿¿¿) and Atomic Orbitals 54
3.5 The Four Quantum Numbers 56
3.6 The Polyelectronic Atoms and the Filling of Orbitals for the Atoms of the Elements 58
3.7 Aufbau Principle 61
3.8 Atomic Properties 62
3.8.1 Orbitals Energies and Shielding 62
3.8.2 Term Symbols: Coupling of Spin and Orbital Angular Momentum 63
3.8.3 Periodic Properties - Atomic Radius 67
3.8.4 Periodic Properties - Ionization Potential (IP) 67
3.8.5 Periodic Properties - Electron Affinity (EA) 71
3.8.6 Periodic Properties - Electronegativity (¿¿¿¿) 74
3.8.7 Periodic Properties - Hardness (¿¿¿¿) 75
References 77
4. Symmetry 79
4.1 Introduction 79
4.2 Symmetry Concepts 79
4.2.1 Symmetry Operation 79
4.2.2 Symmetry Element 79
4.2.3 Symmetry Elements and Operations 80
4.3 Point Groups 84
4.3.1 Special Groups and Platonic Solids/Polyhedra 85
4.3.2 Examples of the Use of the Scheme for Determining Point Groups 88
4.4 Optical Isomerism and Symmetry 92
4.4.1 Dichloro-Allene Derivatives (C3H2Cl2) 92
4.4.2 Tartaric Acid 93
4.4.3 Cylindrical Helix Molecules 93
4.5 Fundamentals of Group Theory 93
4.5.1 C2v Point Group 95
4.5.2 Explanation of the Character Table 96
4.5.3 Generation of the Irreducible Representations (C2v Case) 97
4.5.4 Notation for Irreducible Representations 97
4.5.5 Some Important Properties of the Characters and their Irreducible Representations 98
4.5.6 Nonindependence of x and y Transformations (Higher Order Rotations) 98
4.6 Selected Applications of Group Theory 101
4.6.1 Generation of a Reducible Representation to Describe a Molecule 101
4.6.2 Determining the IR and Raman Activity of Vibrations in Molecules 104
4.6.3 Determining the Vibrational Modes of Methane, CH4 105
4.6.4 Determining the Irreducible Representations and Symmetry of the Central Atom's Atomic Orbitals that Form Bonds 107
4.7 Symmetry Adapted Linear Combination (SALC) of Orbitals 111
4.7.1 Sigma Bonding with Hydrogen as Terminal Atom 111
4.7.2 Sigma and Pi Bonding with Atoms Other than Hydrogen as Terminal Atom 114
Appendix 4.1 Some Additional useful Character Tables 120
References 122
5. Covalent Bonding 123
5.1 Introduction 123
5.1.1 Lewis Structures and the Octet Rule 123
5.1.2 Valence Shell Electron Pair Repulsion Theory (VSEPR) 126
5.2 Valence Bond Theory (VBT) 127
5.2.1 H2 and Valence Bond Theory 129
5.2.2 Ionic Contributions to Covalent Bonding 130
5.2.3 Polyatomic Molecules and Valence Bond Theory 131
5.3 Molecular Orbital Theory (MOT) 132
5.3.1 H2 132
5.3.2 Types of Orbital Overlap 137
5.3.3 Writing Generalized Wave Functions 138
5.3.4 Brief Comments on Computational Methods and Computer Modeling 139
5.3.5 Homonuclear Diatomic Molecules (A2) 140
5.3.6 Heteronuclear Diatomic Molecules and Ions (AB; HX) - Sigma Bonds Only 144
5.3.7 Heteronuclear Diatomic Molecules and Ions (AB) - Sigma and Pi Bonds 147
5.4 Understanding Reactions and Electron Transfer (Frontier Molecular Orbital Theory) 150
5.4.1 Angular Overlap 151
5.4.2 H+ +OH- 151
5.4.3 H2 +D2 152
5.4.4 H2 +F2 153
5.4.5 H2 +C2 154
5.4.6 H2 +N2 (also CO+H2) 154
5.4.7 Dihalogens as Oxidants 156
5.4.8 O2 as an Oxidant and its Reaction with H2S and HS- 157
5.5 Polyatomic Molecules and Ions 161
5.5.1 H3+ Molecular Cation 161
5.5.2 BeH2 - Linear Molecule with Sigma Bonds Only 163
5.5.3 H2O - Angular Molecule with Sigma Bonds Only 165
5.6 Tetrahedral and Pyramidal Species with Sigma Bonds only (CH4, NH4+, SO42-) 168
5.6.1 CH4 168
5.6.2 NH3 (C3v) 170
5.6.3 BH3 and the Methyl Cation, CH3+ (D3h) 172
5.7 Triatomic Compounds and Ions Involving ¿¿¿¿ Bonds (A3, AB2, and ABC) 175
5.7.1 A3 Linear Species 175
5.7.2 AB2 Linear Species CO2 (COS and N2O) 178
5.7.3 O3, NO2-, and SO2 (Angular Molecules) 180
5.8 Planar Species (BF3, NO3-, CO32-, SO3) 182
Appendix 5.1 Bond Energies for Selected Bonds 184
Appendix 5.2 Energies of LUMOs and HOMOs 185
References 186
6. Bonding in Solids 189
6.1 Introduction 189
6.2 Covalent Bonding in Metals: Band Theory 189
6.2.1 Atomic Orbital Combinations for Metals 189
6.2.2 Metal Conductors 191
6.2.3 Semiconductors and Insulators 191
6.2.4 Fermi Level 193
6.2.5 Density of States (DOS) 194
6.2.6 Doping of Semiconductors 195
6.2.7 Structures of Solids 196
6.3 Ionic Solids 200
6.3.1 Solids AX Stoichiometry 200
6.3.2 Solids with Stoichiometry of AX2, AO2, A2O3, ABO3 (Perovskite), AB2O4 (Spinel) 203
6.3.3 Crystal Radii 205
6.3.4 Radius Ratio Rule 205
6.3.5 Lattice Energy 207
6.3.6 Born-Haber Cycle 209
6.3.7 Thermal Stability of Ionic Solids 210
6.3.8 Defect Crystal Structures 212
6.4 Nanoparticles and Molecular Clusters 214
References 217
7. Acids and Bases 219
7.1 Introduction 219
7.2 Arrhenius and Bronsted-Lowry Definitions 219
7.3 Hydrolysis of Metal-Water Complexes 222
7.4 Hydration of Anhydrous Acidic and Basic Oxides 223
7.4.1 Acidic Oxides 223
7.4.2 Basic Oxides 224
7.4.3 Amphoteric Oxides 224
7.5 Solvent System Definition 224
7.5.1 Leveling Effect 225
7.6 Gas Phase Acid-Base Strength 225
7.6.1 H3+ as a Reactant 227
7.7 Lewis Definition 227
7.7.1 MOT 228
7.7.2 Molecular Iodine Adducts or Complexes as Examples 228
7.7.3 Thermodynamics of Lewis Acid-Base Reactions 229
7.7.4 Lewis Acid-Base Reactions of CO2 and I2 with Water and Hydroxide Ion 230
7.7.5 Lewis Acid-Base Competitive Reactions 232
7.8 Classification of Acids and Bases 232
7.8.1 Irving-Williams Stability Relationship for the First Transition Metal Series 232
7.8.2 Class "a" and "b" Acids and Bases 233
7.8.3 Hard Soft Acid Base (HSAB) Theory 233
7.9 Acid-Base Properties of Solids 235
References 235
8. Introduction to Transition Metals 237
8.1 Introduction 237
8.2 Coordination Geometries 237
8.3 Nomenclature 240
8.3.1 Complex Ion is Positive 241
8.3.2 Complex Ion is Negative 242
8.3.3 Complex Ion with Multiple Ligands 242
8.3.4 Complex Ion with Ligand that can Bind with More Than One Atom (Ambidentate) 243
8.3.5 Complex Ion with Multidentate Ligands 243
8.3.6 Two Complex Ions with a Bridging Ligand 243
8.4 Bonding and Isomers for Octahedral Geometry 243
8.4.1 Ionization Isomerism 244
8.4.2 Hydrate (Solvate) Isomers 244
8.4.3 Coordination Isomerism 245
8.4.4 Linkage Isomerism 245
8.4.5 Geometrical Isomerism - Four Coordination 246
8.4.6 Optical Isomerism in Octahedral Geometry 248
8.5 Bonding Theories for Transition Metal Complexes 250
8.5.1 Valence Bond Theory 251
8.5.2 Crystal Field Theory 252
8.6 Molecular Orbital Theory 268
8.6.1 Case 1 - Octahedral Geometry (Sigma Bonding Only) 268
8.6.2 Case 2 - Octahedral Geometry (Sigma Bonding Plus Ligand ¿¿¿¿ Donor) 271
8.6.3 Case 3 - Octahedral Geometry (Sigma Bonding Plus Ligand ¿¿¿¿ Acceptor) 272
8.7 Angular Overlap Model 274
8.7.1 AOM and ¿¿¿¿ Ligand Donor Bonding 277
8.7.2 AOM and ¿¿¿¿ Ligand Acceptor Bonding 278
8.7.3 MOT, Electrochemistry, and the Occupancy of Electrons in d Orbitals in Oh 278
8.7.4 AOM and Other Geometries 279
8.8 More on Spectroscopy of Metal-Ligand Complexes 281
8.8.1 Charge Transfer Electronic Transitions 282
8.8.2 Electronic Spectra, Spectroscopic Terms, and the Energies of the Terms for d¿d Transitions 283
8.8.3 Energy and Spatial Description of the Electron Transitions Between t2g and eg * Orbitals 296
8.8.4 More Details on Correlation Diagrams 297
8.8.5 Luminescence 299
8.8.6 Magnetism and Spin Crossover in Octahedral Complexes and Natural Minerals 301
8.8.7 Note about f Orbitals in Cubic Symmetry (Oh) 303
References 303
9. Reactivity of Transition Metal Complexes: Thermodynamics, Kinetics and Catalysis 305
9.1 Thermodynamics Introduction 305
9.1.1 Successive Stability Constants on Water Substitution 305
9.1.2 The Chelate Effect 307
9.2 Kinetics of Ligand Substitution Reactions 308
9.2.1 Kinetics of Water Exchange for Aqua Complexes 310
9.2.2 Intimate Mechanisms for Ligand Substitution Reactions 310
9.2.3 Kinetic Model and Activation Parameters 311
9.2.4 Dissociative Versus Associative Preference for Octahedral Ligand Substitution Reactions 314
9.2.5 Stoichiometric Mechanisms 315
9.2.6 Tests for Reaction Mechanisms 320
9.3 Substitution in Octahedral Complexes 321
9.3.1 Examples of Dissociative Activated Mechanisms 321
9.3.2 Associative Activated Mechanisms 322
9.4 Intimate Mechanisms Affected by Steric Factors (Dissociative Preference) 324
9.4.1 Intimate Mechanisms Affected by Ligands in Cis versus Trans Positions (Dissociative Preference) 324
9.4.2 Base Hydrolysis 325
9.5 Intimate versus Stoichiometric Mechanisms 327
9.6 Substitution in Square Planar Complexes (Associative Activation Predominates) 328
9.6.1 Effect of Leaving Group 330
9.6.2 Effect of Charge 330
9.6.3 Nature of the Intermediate - Electronic Factors 330
9.6.4 Nature of the Intermediate - Steric Factors 331
9.7 Metal Electron Transfer Reactions 332
9.7.1 Outer Sphere Electron Transfer 333
9.7.2 Cross Reactions 337
9.7.3 Inner Sphere Electron Transfer 339
9.8 Photochemistry 341
9.8.1 Redox 341
9.8.2 Photosubstitution Reactions d¿d 341
9.8.3 LMCT and Photoreduction 342
9.8.4 MLCT Simultaneous Substitution and Photo-Oxidation Redox 342
9.9 Effective Atomic Number (EAN) Rule or the Rule of 18 342
9.10 Thermodynamics and Kinetics of Organometallic Compounds 344
9.11 Electron Transfer to Molecules during Transition Metal Catalysis 345
9.12 Oxidation Addition (OXAD) and Reductive Elimination (Redel) Reactions 346
9.13 Metal Catalysis 347
9.13.1 OXO or Hydroformylation Process 348
9.13.2 Heck Reaction 350
9.13.3 Methyl Transferases 350
9.13.4 Examples of Abiotic Organic Synthesis (Laboratory and Nature) 351
9.13.5 The Haber Process Revisited 353
References 353
10. Transition Metals in Natural Systems 356
10.1 Introduction 356
10.2 Factors Governing Metal Speciation in the Environment and in Organisms 356
10.3 Transition Metals Essential for Life 358
10.4 Important Environmental Iron and Manganese Reactions 359
10.4.1 Oxidation of Fe2+ and Mn2+ by O2 - Environmentally Important Metal Electron Transfer Reactions 360
10.4.2 Redox Properties of Iron-Ligand Complexes 363
10.4.3 Metal Ions Exhibiting Outer Sphere Electron Transfer 364
10.5 Oxygen (O2) Storage and Transport 364
10.5.1 Hemoglobin 365
10.5.2 Hemocyanin and Hemerythrin 368
10.6 Oxidation of CH4, Hydrocarbons, NH4+ 368
10.6.1 Cytochrome P450: An Example of Cytochrome (Heme - O2) Redox Chemistry 369
10.6.2 Conversion of NH4+ to NO3- (Nitrification or Aerobic Ammonium Oxidation) 371
10.7 Oxygen Production in Photosynthesis 372
References 374
11. Solid Phase Iron and Manganese Oxidants and Reductants 377
11.1 Introduction 377
11.2 Reduction of Solid MnO2 and Fe(OH)3 by Sulfide 377
11.2.1 Fe(III) and Mn(IV) Electron Configurations 378
11.2.2 MnO2 Reaction with Sulfide 379
11.2.3 Fe(OH)3 Reaction with Sulfide 382
11.3 Pyrite, FeS2, Oxidation 384
11.3.1 Pyrite Reacting with O2 384
11.3.2 Pyrite Reacting with Soluble Fe(III) 385
11.3.3 Pyrite Reacting with Dihalogens and Cr2+ 387
References 388
12. Metal Sulfides in the Environment and in Bioinorganic Chemistry 390
12.1 Introduction 390
12.2 Idealized Molecular Reaction Schemes from Soluble Complexes to ZnS and CuS Solids 391
12.3 Nanoparticle Size and Filtration 394
12.4 Ostwald Ripening versus Oriented Attachment 394
12.5 Metal Availability and Detoxification for MS Species 396
12.6 Iron Sulfide Chemistry 396
12.6.1 FeSmack (Mackinawite) 396
12.6.2 FeSmack Conversion to Pyrite, FeS2 397
12.6.3 FeS as a Catalyst in Organic Compound Formation 400
12.6.4 FeS as an Electron Transfer Agent in Biochemistry 400
12.7 More on the Nitrogen Cycle (Nitrate Reduction, Denitrification, and Anammox) 402
Appendix 12.1 PbS Nanoparticle Model and Size Ranges of Natural Materials 404
References 404
13. Kinetics and Thermodynamics of Metal Uptake by Organisms 406
13.1 Introduction 406
13.1.1 Conditional Metal-Ligand Stability Constants 407
13.1.2 Thermodynamic Metal-Ligand Stability Constants 409
13.2 Metal Uptake Pathways 410
13.2.1 Ion Channels for Potassium 411
13.2.2 Metal Uptake by Cells via Ligands on Membranes 413
13.2.3 Evaluation of kf , kd, and KcondM'L' from Laboratory and Natural Samples 418
References 420
Index 421
Understanding the atomic structure of the atom is important to understanding how atoms combine to build molecules including minerals, aqueous materials, and gases. However, nature builds the atoms of the elements starting with the simplest elements, hydrogen and helium, which are the most abundant in the cosmos. As described in the Big Bang Theory [1-6 and references therein], the origin of the universe (and the periodic table) started with elementary particles under enormous gravitational attraction concentrated in an extremely densely packed point, which exploded causing an expanding universe and the release of enormous energy and fundamental particles. Within 0.01 s of the Big Bang, temperatures have been predicted to be in the range of -. After 100 s and on cooling to , the elementary particles began to combine under the force of gravity. Here, positively charged protons and neutral neutrons start to combine to form the lighter elements (nucleosynthesis) and their isotopes. Their combination in the nucleus occurs by the strong force, which is the short-range attractive force between protons and neutrons that binds these particles in the nucleus while overcoming the repulsive force of the protons with each other. At this time and under these conditions, the electrons are totally ionized from the nucleus and cannot combine with the elements until cooling occurs at about . At this lower temperature, the electromagnetic force begins to take effect and the combination of the electrons with the positive nuclei to form neutral atoms occurs. Once there is a buildup of neutral atoms, chemical processes can occur that eventually lead to life and biological processes.
To understand the energies associated with a wide range of processes at temperatures from absolute zero to these extreme Big Bang temperatures, the Boltzmann energy-temperature relationship (Equation 1.1) provides perspective:
where ; multiplying by (as ) gives E in units of eV (electron volts). Figure 1.1 is a plot of E (eV) versus that gives the temperature and corresponding energy at which several well-known processes occur. Multiplying by Avogadro's number provides R, the gas constant, and E in units of .
Figure 1.1 Log-log plot of energy versus temperature; circles include the temperature at which some familiar chemical and physical processes including hydrothermal vents (~360 °C) found on deep ocean ridges occur. Triangles indicate the T and E parameters for nucleosynthesis in a sun that has 10-20 times the mass of our sun
Before continuing with the process of nucleosynthesis, it is necessary to define the general symbol used for nuclides, which includes their nuclear and charged properties, as where El is the element symbol, (total number of protons and neutrons or total nucleons), (number of protons) and is the charge due to loss or gain of electrons. The difference of A - Z equals the number of neutrons (N). Isotopes of an element have different atomic masses and the same atomic number due to a different number of neutrons in the nucleus.
Immediately after the Big Bang, the buildup of He (and other light elements) from protons and neutrons occurred through several multistep nuclear processes. Equations 1.2-1.5 show one example (positive charges are omitted for simplicity after Equation 1.2). The free neutron has a half-life of 13 min so the formation of hydrogen (Equation 1.2) occurs rapidly with formation of the electron ( or ) and one of the neutrinos (another radiation component; see Table 1.1). The first nuclear reaction in this sequence is between the proton () and the neutron to form positively charged deuterium (deuteron, ), and the buildup of eventually leads to positively charged tritium () and positively charged helium () formation. For example, continued reaction of the proton with the deuteron produces the doubly charged (Helium-3). Under these extreme temperatures , the repulsive forces of the positively charged particles can be overcome so that the charged particles combine in the nucleus, which has a size on the order of diameter. (At higher temperatures, the can decay due to photodissociation.) can then combine with another neutron to form (also known as the alpha particle; Helium-4) where indicates gamma rays that are at the high energy region of the electromagnetic spectrum (Figure 1.2). The energy released for Equation 1.3 and subsequent reactions is substantial, and maintains or increases the initial temperature. Because of these extreme temperatures, the elements were actually in a plasma state.
Table 1.1 Some important atomic and subatomic particles. One atomic mass unit (amu) equals 1.6606 × 10- 27 kg (the atomic mass constant) and the elementary charge is 1.602 × 10- 19 coulomb (C). Spin is in units of h/(2p) or (h = Planck's constant). ß- and ß+ are ejected from the nucleus with ß+ formally an antiparticle to ß-; the reaction of ß- with ß+ leads to their annihilation and release of ?-ray energy. The electron neutrino, ?e or ?-, and the positron neutrino, e or ?+ (also known as the antineutrino), account for excess energy release during nuclear reactions (there are two other neutrinos and antineutrinos that are not important to this discussion)
Figure 1.2 The electromagnetic spectrum is given in terms of wavelength and frequency (, bottom axis) and energy in eV and Joule per atom (top axis; see (Equation 3.2; recall ). . The types of spectroscopic techniques for these energy regions are at the bottom
The proton and the neutron interconvert in atoms to increase nuclear stability via the weak nuclear force that operates at [4]. Equation 1.2 is an example of spontaneous beta emission (decay) of the neutron. Other examples of neutron conversion to a proton via reaction with a particle and an electron neutrino are given in Equations 1.6a and 1.6b, respectively. The proton is stable to decay so requires another particle (Equation 1.6c) or energy (antineutrino; Equation 1.6d) to transform it to a neutron.
Further buildup of the elements occurs in the stars and during stellar explosions (supernovae), and the energy released for those reactions above and subsequent reactions is substantial (Section 1.4). The initial process in stellar nuclear synthesis is hydrogen burning (our sun is undergoing this process), which is described with the following set of equations (1.7a-1.7d) and which occurs at temperatures of .
is the positron (also given as ) and is one of the neutrinos.
The sum of Equations 1.7a-1.7c gives the overall nuclear reaction to produce the alpha particle, (Equation 1.7d), and is one of three proton-proton chain sequences [6].
Once helium is produced, the core of the star continues to collapse and gravitational forces drive the temperature up to where fuses with itself in a reaction sequence (Equations 1.8a-1.8d) to form (half-life of ), which rapidly fuses with another to form carbon (the triple process), which can react further with helium. This sequence is one of the helium burning reaction sequences.
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