(BQ) Part 1 book Descriptive inorganic chemistry has contents: The electronic structure of the atom A review; an overview of the periodic table; covalent bondin; ionic bonding; solvent systems and acid base behavior; oxidation and reduction,...and other contents.
Trang 4Acquisitions Editors: Jessica Fiorillo/Kathryn Treadway
Marketing Director: John Britch
Media Editor: Dave Quinn
Cover and Text Designer: Vicki Tomaselli
Senior Project Editor: Mary Louise Byrd
Illustrations: Network Graphics/Aptara
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Production Coordinator: Susan Wein
Composition: Aptara
Printing and Binding: World Color Versailles
Library of Congress Control Number: 2009932448
Trang 5CHAPTER 1 The Electronic Structure of the Atom: A Review 1
CHAPTER 2 An Overview of the Periodic Table 19
CHAPTER 7 Solvent Systems and Acid-Base Behavior 137
CHAPTER 11 The Group 1 Elements: The Alkali Metals 245
CHAPTER 12 The Group 2 Elements: The Alkaline Earth Metals 271
CHAPTER 15 The Group 15 Elements: The Pnictogens 363
CHAPTER 16 The Group 16 Elements: The Chalcogens 409
CHAPTER 17 The Group 17 Elements: The Halogens 453
CHAPTER 18 The Group 18 Elements: The Noble Gases 487
CHAPTER 19 Transition Metal Complexes 499
CHAPTER 20 Properties of the 3d Transition Metals 533
CHAPTER 21 Properties of the 4d and 5d Transition Metals 579
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CHAPTER 24 The Rare Earth and Actinoid Elements 651w
Overview
iii
Trang 7Atomic Absorption Spectroscopy 2
1.1 The Schrödinger Wave Equation and Its
1.2 Shapes of the Atomic Orbitals 5
1.4 Ion Electron Confi gurations 14
1.5 Magnetic Properties of Atoms 15
1.6 Medicinal Inorganic Chemistry:
CHAPTER 2
2.1 Organization of the Modern
2.2 Existence of the Elements 23
2.3 Stability of the Elements and Their Isotopes 24
The Origin of the Shell Model of the Nucleus 26
2.4 Classifi cations of the Elements 27
2.5 Periodic Properties: Atomic Radius 29
2.6 Periodic Properties: Ionization Energy 33
2.7 Periodic Properties: Electron Affi nity 35
CHAPTER 3
3.1 Models of Covalent Bonding 42
3.2 Introduction to Molecular Orbitals 43
3.3 Molecular Orbitals for Period 1
3.6 A Brief Review of Lewis Structures 51
3.9 Valence-Shell Electron-Pair Repulsion Rules 543.10 The Valence-Bond Concept 593.11 Network Covalent Substances 61
Concrete: An Old Material with a New Future 109
v
Trang 8CHAPTER 6
6.1 Thermodynamics of the Formation
6.2 Formation of Ionic Compounds 120
6.4 Thermodynamics of the Solution
Process for Ionic Compounds 124
6.5 Formation of Covalent Compounds 127
6.6 Thermodynamic versus Kinetic Factors 129
Cyanide and Tropical Fish 148
7.4 Trends in Acid-Base Behavior 148
Superacids and Superbases 150
7.5 Acid-Base Reactions of Oxides 153
7.7 Pearson Hard-Soft Acid-Base Concepts 156
7.8 Applications of the HSAB Concept 158
CHAPTER 8
8.3 Determination of Oxidation Numbers
8.4 The Difference between Oxidation
8.5 Periodic Variations of Oxidation
Numbers 172
Chemosynthesis: Redox Chemistry on the
8.7 Quantitative Aspects of Half-Reactions 176
8.8 Electrode Potentials as Thermodynamic Functions 1778.9 Latimer (Reduction Potential) Diagrams 1788.10 Frost (Oxidation State) Diagrams 180
Trang 9vii
10.7 Biological Aspects of Hydrogen
Bonding 241
Is There Life Elsewhere in Our Solar System? 242
10.8 Element Reaction Flowchart 242
CHAPTER 11
The Group 1 Elements: The Alkali
Metals 245
11.2 Features of Alkali Metal Compounds 247
11.3 Solubility of Alkali Metal Salts 249
Trang 1015.2 Contrasts in the Chemistry
of Nitrogen and Phosphorus 365
15.3 Overview of Nitrogen Chemistry 368
The First Dinitrogen Compound 369
15.10 Nitrous Acid and Nitrites 385
15.11 Nitric Acid and Nitrates 386
15.12 Overview of Phosphorus Chemistry 389
Paul Erhlich and His “Magic Bullet” 401
15.29 Element Reaction Flowcharts 402
CHAPTER 16
The Group 16 Elements:
16.2 Contrasts in the Chemistry of
Oxygen Isotopes in Geology 41216.4 Bonding in Covalent Oxygen
Compounds 41816.5 Trends in Oxide Properties 419
16.11 Overview of Sulfur Chemistry 426
17.2 Contrasts in the Chemistry of
The Fluoridation of Water 45917.4 Hydrogen Fluoride and Hydrofl uoric Acid 46017.5 Overview of Chlorine Chemistry 462
Trang 1117.10 Chlorine Oxyacids and Oxyanions 471
The Discovery of the Perbromate Ion 474
17.11 Interhalogen Compounds and
The Group 18 Elements:
18.2 Unique Features of Helium 489
18.3 Uses of the Noble Gases 489
18.4 A Brief History of Noble Gas
Compounds 491
Is It Possible to Make Compounds of the
Platinum Complexes and Cancer Treatment 506
19.5 Naming Transition Metal Complexes 507
19.6 An Overview of Bonding Theories
of Transition Metal Compounds 510
19.8 Successes of Crystal Field Theory 517
The Earth and Crystal Structures 52119.9 More on Electronic Spectra 521
19.11 Thermodynamic versus Kinetic Factors 52519.12 Synthesis of Coordination Compounds 52619.13 Coordination Complexes and the
Radiopharmaceutical 58821.7 The Platinum Group Metals 58921.8 Group 8: Ruthenium and Osmium 59021.9 Group 9: Rhodium and Iridium 59121.10 Group 10: Palladium and Platinum 59121.11 Group 11: Silver and Gold 591
Contents
Trang 1222.5 Element Reaction Flowchart 608
23.4 Solvents for Organometallic Chemistry 614
23.5 Main Group Organometallic
Compounds 615
The Death of Karen Wetterhahn 623
23.6 Organometallic Compounds of the
23.7 Transition Metal Carbonyls 625
23.8 Synthesis and Properties of Simple
23.9 Reactions of Transition Metal Carbonyls 632
23.10 Other Carbonyl Compounds 633
23.11 Complexes with Phosphine Ligands 634
23.12 Complexes with Alkyl, Alkene, and
23.15 Complexes with 6-Arene Ligands 642
23.16 Complexes with Cycloheptatriene and
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The Rare Earth and Actinoid Elements 651w
Appendix 1 Thermodynamic Properties of
Some Selected Inorganic Compounds A-1Appendix 2 Charge Densities of Selected
Ions A-13Appendix 3 Selected Bond Energies A-16Appendix 4 Ionization Energies of Selected
Metals A-18Appendix 5 Electron Affi nities of Selected
Nonmetals A-20Appendix 6 Selected Lattice Energies A-21Appendix 7 Selected Hydration Enthalpies A-22Appendix 8 Selected Ionic Radii A-23
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Appendix 9 Standard Half-Cell Electrode
Potentials of Selected Elements A-25w
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Appendix 10 Electron Confi guration
Trang 13What Is Descriptive Inorganic Chemistry?
Descriptive inorganic chemistry was traditionally concerned with the
prop-erties of the elements and their compounds Now, in the renaissance of
the subject, with the synthesis of new and novel materials, the properties are
being linked with explanations for the formulas and structures of compounds
together with an understanding of the chemical reactions they undergo In
addition, we are no longer looking at inorganic chemistry as an isolated subject
but as a part of essential scientifi c knowledge with applications throughout
science and our lives Because of a need for greater contextualization, we have
added more features and more applications
In many colleges and universities, descriptive inorganic chemistry is offered
as a sophomore or junior course In this way, students come to know something
of the fundamental properties of important and interesting elements and their
compounds Such knowledge is important for careers not only in pure or applied
chemistry but also in pharmacy, medicine, geology, and environmental science
This course can then be followed by a junior or senior course that focuses on
the theoretical principles and the use of spectroscopy to a greater depth than
is covered in a descriptive text In fact, the theoretical course builds nicely on
the descriptive background Without the descriptive grounding, however, the
theory becomes sterile, uninteresting, and irrelevant
Education has often been a case of the “swinging pendulum,” and this
has been true of inorganic chemistry Up until the 1960s, it was very much
pure descriptive, requiring exclusively memorization In the 1970s and 1980s,
upper-level texts focused exclusively on the theoretical principles Now it is
ap-parent that descriptive is very important—not the traditional memorization of
facts but the linking of facts, where possible, to underlying principles Students
need to have modern descriptive inorganic chemistry as part of their
educa-tion Thus, we must ensure that chemists are aware of the “new descriptive
inorganic chemistry.”
xi
Trang 15Inorganic chemistry goes beyond academic interest: it is an
im-portant part of our lives.
Inorganic chemistry is interesting—more than that—it is exciting! So much
of our twenty-fi rst-century science and technology rely on natural and
syn-thetic materials, often inorganic compounds, many of which are new and novel
Inorganic chemistry is ubiquitous in our daily lives: household products, some
pharmaceuticals, our transportation—both the vehicles themselves and the
synthesis of the fuels—battery technology, and medical treatments There is
the industrial aspect, the production of all the chemicals required to drive our
economy, everything from steel to sulfuric acid to glass and cement
Environ-mental chemistry is largely a question of the inorganic chemistry of the
atmo-sphere, water, and soil Finally, there are the profound issues of the inorganic
chemistry of our planet, the solar system, and the universe
This textbook is designed to focus on the properties of selected interesting,
important, and unusual elements and compounds However, to understand
inorganic chemistry, it is crucial to tie this knowledge to the underlying
chemi-cal principles and hence provide explanations for the existence and behavior
of compounds For this reason, almost half the chapters survey the relevant
concepts of atomic theory, bonding, intermolecular forces, thermodynamics,
acid-base behavior, and reduction-oxidation properties as a prelude to, and
preparation for, the descriptive material
For this fi fth edition, the greatest change has been the expansion of coverage
of the 4d and 5d transition metals to a whole chapter.
The heavier transition metals have unique trends and patterns, and the
new chapter highlights these Having an additional chapter on transition
met-als met-also better balances the coverage between the main group elements and the
transition elements.
Also, the fi fth edition has a second color With the addition of a second color,
fi gures are much easier to understand, and tables and text are easier to read.
On a chapter-by-chapter basis, the signifi cant improvements are as follows:
Chapter 1: The Electronic Structure of the Atom: A Review
The Introduction and Section 1.3, The Polyelectronic Atom, have been revised
Chapter 3: Covalent Bonding
Section 3.11, Network Covalent Substances, has a new subsection: Amorphous
Silicon
Chapter 4: Metallic Bonding
Section 4.6, Nanometal Particles, was added
Section 4.7, Magnetic Properties of Metals, was added
xiii
Trang 16Chapter 5: Ionic Bonding
Section 5.3, Polarization and Covalency, has a new subsection: The Covalent Boundary
Ionic-Section 5.4, Ionic Crystal Structures, has a new subsection: Quantum Dots
Chapter 9: Periodic Trends
Section 9.3, Isoelectronic Series in Covalent Compounds, has been revised and improved
Section 9.8, The “Knight’s Move” Relationship, has been revised and improved
Chapter 10: Hydrogen
Section 10.4, Hydrides, has a revised and expanded subsection: Ionic Hydrides
Chapter 11: The Group 1 Elements
Section 11.14, Ammonium Ion as a Pseudo–Alkali-Metal Ion, moved from Chapter 9
Chapter 13: The Group 13 Elements
Section 13.10, Aluminides, was added
Chapter 14: The Group 14 Elements
Section 14.2, Contrasts in the Chemistry of Carbon and Silicon, was added.Section 14.3, Carbon, has a new subsection: Graphene
Section 14.7, Carbon Dioxide, has a new subsection: Carbonia
Chapter 15: The Group 15 Elements
Section 15.2, Contrasts in the Chemistry of Nitrogen and Phosphorus, was added
Section 15.18, The Pnictides, was added
Chapter 16: The Group 16 Elements
Section 16.2, Contrasts in the Chemistry of Oxygen and Sulfur, was added.Section 16.14, Sulfi des, has a new subsection: Disulfi des
Chapter 17: The Group 17 Elements
Section 17.2, Contrasts in the Chemistry of Fluorine and Chlorine, was added.Section 17.12, Cyanide Ion as a Pseudo-halide Ion, moved from Chapter 9
Chapter 18: The Group 18 Elements
Section 18.7, Other Noble Gas Compounds, was added
Chapter 19: Transition Metal Complexes
Section 19.10, Ligand Field Theory, was added
Chapter 20: Properties of the 3d Transition Metals
Section 20.1, Overview of the 3d Transition Metals, was added
Chapter 21: Properties of the 4d and 5d Transition Metals
N EW C HAPTERadded (for details, see the previous page)
Chapter 24: The Rare Earth and Actinoid Elements
This chapter has been signifi cantly revised with the new subsections Scandium, Yttrium, and Thorium
Trang 17xv
ALSO
Video Clips
Descriptive inorganic chemistry by defi nition is visual, so what better way to
appreciate a chemical reaction than to make it visual? We now have a series of
at least 60 Web-based video clips to bring some of the reactions to life The text
has a margin icon to indicate where a reaction is illustrated
Text Figures and Tables
All the illustrations and tables in the book are available as jpg fi les for inclusion
in PowerPoint presentations on the instructor side of the Web site at
www.whfreeman.com/descriptive5e
Additional Resources
A list of relevant S CIENTIFIC A MERICANarticles is found on the text Web site
at www.whfreeman.com/descriptive5e The text has a margin icon to indicate
where a Scientifi c American article is available
Supplements
The Student Solutions Manual, ISBN: 1-4292-2434-7 contains the worked
solutions to all the odd-numbered end-of-chapter problems
The Companion Web Site www.whfreeman.com/descriptive5e
Contains the following student-friendly materials: Chapter 24: The Rare Earth
and Actinoid Elements, Appendices, Lab Experiments, Tables, and over 50 useful
videos of elements and metals in reactions and oxidations
Instructor’s Resource CD-ROM, ISBN: 1-4292-2428-2
Includes PowerPoint and videos as well as all text art and solutions to all
prob-lems in the book
This textbook was written to pass on to another generation our fascination
with descriptive inorganic chemistry Thus, the comments of readers, both
stu-dents and instructors, will be sincerely appreciated Any suggestions for added
or updated additional readings are also welcome Our current e-mail addresses
aregrcanham@swgc.mun.ca and T.L.Overton@hull.ac.uk
Preface
Trang 19Many thanks must go to the team at W H Freeman and Company who have
contributed their talents to the fi ve editions of this book We offer our sincere
gratitude to the editors of the fi fth edition, Jessica Fiorillo, Kathryn Treadway, and
Mary Louise Byrd; of the fourth edition, Jessica Fiorillo, Jenness Crawford, and
Mary Louise Byrd; of the third edition, Jessica Fiorillo and Guy Copes; of the
second edition, Michelle Julet and Mary Louise Byrd; and a special thanks to
Deborah Allen, who bravely commissioned the fi rst edition of the text Each one of
our fabulous editors has been a source of encouragement, support, and helpfulness
We wish to acknowledge the following reviewers of this edition, whose
criticisms and comments were much appreciated: Theodore Betley at Harvard
University; Dean Campbell at Bradley University; Maria Contel at Brooklyn
College (CUNY); Gerry Davidson at St Francis College; Maria Derosa at
Carleton University; Stan Duraj at Cleveland State University; Dmitri Giarkios
at Nova Southeastern University; Michael Jensen at Ohio University–Main
Campus; David Marx at the University of Scranton; Joshua Moore at Tennessee
State University–Nashville; Stacy O’Reilly at Butler University; William
Pen-nington at Clemson University; Daniel Rabinovich at the University of North
Carolina at Charlotte; Hal Rogers at California State University–Fullerton;
Thomas Schmedake at the University of North Carolina at Charlotte; Bradley
Smucker at Austin College; Sabrina Sobel at Hofstra University; Ronald
Strange at Fairleigh Dickinson University–Madison; Mark Walters at New
York University; Yixuan Wang at Albany State University; and Juchao Yan at
Eastern New Mexico University; together with prereviewers: Londa Borer at
California State University–Sacramento; Joe Fritsch at Pepperdine
Univer-sity; Rebecca Roesner at Illinois Wesleyan University, and Carmen Works at
Sonoma College
We acknowledge with thanks the contributions of the reviewers of the
fourth edition: Rachel Narehood Austin at Bates College; Leo A Bares at the
University of North Carolina—Asheville; Karen S Brewer at Hamilton College;
Robert M Burns at Alma College; Do Chang at Averett University; Georges
Dénès at Concordia University; Daniel R Derringer at Hollins University;
Carl P Fictorie at Dordt College; Margaret Kastner at Bucknell University;
Michael Laing at the University of Natal, Durban; Richard H Langley at
Stephen F Austin State University; Mark R McClure at the University of North
Carolina at Pembroke; Louis Mercier at Laurentian University; G Merga at
Andrews University; Stacy O’Reilly at Butler University; Larry D Pedersen
at College Misercordia; Robert D Pike at the College of William and Mary;
William Quintana at New Mexico State University; David F Rieck at Salisbury
University; John Selegue at the University of Kentucky; Melissa M Strait at
Alma College; Daniel J Williams at Kennesaw State University; Juchao Yan at
Eastern New Mexico University; and Arden P Zipp at the State University of
New York at Cortland
xvii
Trang 20And the contributions of the reviewers of the third edition: François Caron
at Laurentian University; Thomas D Getman at Northern Michigan sity; Janet R Morrow at the State University of New York at Buffalo; Robert
Univer-D Pike at the College of William and Mary; Michael B Wells at Cambell versity; and particularly Joe Takats of the University of Alberta for his compre-hensive critique of the second edition
Uni-And the contributions of the reviewers of the second edition: F C Hentz
at North Carolina State University; Michael D Johnson at New Mexico State University; Richard B Kaner at the University of California, Los Angeles; Richard H Langley at Stephen F Austin State University; James M Mayer
at the University of Washington; Jon Melton at Messiah College; Joseph S Merola at Virginia Technical Institute; David Phillips at Wabash College; John
R Pladziewicz at the University of Wisconsin, Eau Claire; Daniel Rabinovich
at the University of North Carolina at Charlotte; David F Reich at Salisbury State University; Todd K Trout at Mercyhurst College; Steve Watton at the Virginia Commonwealth University; and John S Wood at the University of Massachusetts, Amherst
Likewise, the reviewers of the fi rst edition: E Joseph Billo at Boston lege; David Finster at Wittenberg University; Stephen J Hawkes at Oregon State University; Martin Hocking at the University of Victoria; Vake Marganian
Col-at BridgewCol-ater StCol-ate College; Edward Mottel Col-at the Rose-Hulman Institute of Technology; and Alex Whitla at Mount Allison University
As a personal acknowledgment, Geoff Rayner-Canham wishes to
especial-ly thank three teachers and mentors who had a major infl uence on his career: Briant Bourne, Harvey Grammar School; Margaret Goodgame, Imperial Col-lege, London University; and Derek Sutton, Simon Fraser University And he expresses his eternal gratitude to his spouse, Marelene, for her support and encouragement
Tina Overton would like to thank her colleague Phil King for his invaluable suggestions for improvements and his assistance with the illustrations Thanks must also go to her family, Dave, John, and Lucy, for their patience during the months when this project fi lled all her waking hours
Trang 21Chemistry is a human endeavor New discoveries are the result of the work
of enthusiastic people and groups of people who want to explore the
molecular world We hope that you, the reader, will come to share our own
fascination with inorganic chemistry We have chosen to dedicate this book to
two scientists who, for very different reasons, never did receive the ultimate
accolade of a Nobel Prize
Henry Moseley (1887–1915)
Although Mendeleev is identifi ed as the discoverer of the
peri-odic table, his version was based on an increase in atomic mass
In some cases, the order of elements had to be reversed to match
properties with location It was a British scientist, Henry Moseley,
who put the periodic table on a much fi rmer footing by
discov-ering that, on bombardment with electrons, each element
emit-ted X-rays of characteristic wavelengths The wavelengths fi temit-ted
a formula related by an integer number unique to each element
We know that number to be the number of protons With the
es-tablishment of the atomic number of an element, chemists at last
knew the fundamental organization of the table Sadly, Moseley
was killed at the battle of Gallipoli in World War I Thus, one of
the brightest scientifi c talents of the twentieth century died at the
age of 27 The famous American scientist Robert Milliken commented: “Had
the European War had no other result than the snuffi ng out of this young life,
that alone would make it one of the most hideous and most irreparable crimes
in history.” Unfortunately, Nobel Prizes are only awarded to living scientists
In 1924, the discovery of element 43 was claimed, and it was named
mose-leyum; however, the claim was disproved by the very method that Moseley
had pioneered
xix
Trang 22Lise Meitner (1878–1968)
In the 1930s, scientists were bombarding atoms of heavy elements such as uranium with subatomic particles to try to make new ele-ments and extend the periodic table Austrian scientist Lise Meit-ner had shared leadership with Otto Hahn of the German research team working on the synthesis of new elements; the team thought they had discovered nine new elements Shortly after the claimed discovery, Meitner was forced to fl ee Germany because of her Jewish ancestry, and she settled in Sweden Hahn reported to her that one of the new elements behaved chemically just like barium During a famous “walk in the snow” with her nephew, physicist Otto Frisch, Meitner realized that an atomic nucleus could break
in two just like a drop of water No wonder the element formed behaved like barium: it was barium! Thus was born the concept of nuclear fi ssion She informed Hahn of her proposal When Hahn wrote the research paper on the work, he barely mentioned the vital contribution of Meitner and Frisch As a result, Hahn and his colleague Fritz Strassmann received the Nobel Prize Meitner’s fl ash of genius was ignored Only recently has Meitner received the acclaim she deserved by the naming of
an element after her, element 109, meitnerium
Additional reading
Heibron, J L H G J Moseley University of California Press, Berkeley, 1974.Rayner-Canham, M F., and G W Rayner-Canham Women in Chemistry: Their Changing Roles from Alchemical Times to the Mid-Twentieth Century Chemical Heritage Foundation, Philadelphia, 1998
Sime, R L Lise Meitner: A Life in Physics University of California Press, Berkeley, 1996
Weeks, M E., and H M Leicester Discovery of the Elements, 7th ed Journal
of Chemical Education, Easton, PA, 1968
Trang 23Isaac Newton was the original model for the absentminded professor
Supposedly, he always timed the boiled egg he ate at breakfast; one
morning, his maid found him standing by the pot of boiling water,
hold-ing an egg in his hand and gazhold-ing intently at the watch in the bottom
of the pot! Nevertheless, it was Newton who initiated the study of the
electronic structure of the atom in about 1700, when he noticed that
the passage of sunlight through a prism produced a continuous visible
spectrum Much later, in 1860, Robert Bunsen (of burner fame)
inves-tigated the light emissions from fl ames and gases Bunsen observed
that the emission spectra, rather than being continuous, were series of
colored lines (line spectra)
The proposal that electrons existed in concentric shells had its origin
in the research of two overlooked pioneers: Johann Jakob Balmer, a
Swiss mathematician, and Johannes Robert Rydberg, a Swedish physicist
After an undistinguished career in mathematics, in 1885, at the age of
60, Balmer studied the visible emission lines of the hydrogen atom and
found that there was a mathematical relationship between the
wave-lengths Following from Balmer’s work, in 1888, Rydberg deduced a
more general relationship:
where l is the wavelength of the emission line, R H is a constant, later
known as the Rydberg constant, and n f and n i are integers For the
visible lines seen by Balmer and Rydberg, n f had a value of 2 The
Rydberg formula received further support in 1906, when Theodore
Lyman found a series of lines in the far-ultraviolet spectrum of hydrogen,
1.1 The Schrödinger Wave Equation and Its Signifi cance Atomic Absorption Spectroscopy 1.2 Shapes of the Atomic Orbitals 1.3 The Polyelectronic Atom 1.4 Ion Electron Confi gurations 1.5 Magnetic Properties of Atoms 1.6 Medicinal Inorganic Chemistry:
An Introduction
The Electronic Structure
of the Atom: A Review
C HA P TE R 1
1
To understand the behavior of inorganic compounds, we need to study
the nature of chemical bonding Bonding, in turn, relates to the behavior
of electrons in the constituent atoms Our study of inorganic chemistry,
therefore, starts with a review of the models of the atom and a survey of
the probability model’s applications to the electron confi gurations of atoms
and ions.
Trang 24corresponding to the Rydberg formula with n f5 1 Then in 1908, Friedrich Paschen discovered a series of far-infrared hydrogen lines, fi tting the equation
with n f5 3
In 1913, Niels Bohr, a Danish physicist, became aware of Balmer’s and Rydberg’s experimental work and of the Rydberg formula At that time, he was trying to combine Ernest Rutherford’s planetary model for electrons in an atom with Max Planck’s quantum theory of energy exchanges Bohr contended that an electron orbiting an atomic nucleus could only do so at certain fi xed distances and that whenever the electron moved from a higher to a lower orbit, the atom emitted characteristic electromagnetic radiation
Rydberg had deduced his equation from experimental observations of atomic hydrogen emission spectra Bohr was able to derive the same equation from quantum theory, showing that his theoretical work meshed with reality From this result, the Rutherford-Bohr model of the atom of concentric elec-tron “shells” was devised, mirroring the recurring patterns in the periodic table
of the elements (Figure 1.1) Thus the whole concept of electron energy levels can be traced back to Rydberg In recognition of Rydberg’s contribution, excited
atoms with very high values of the principal quantum number, n, are called
Rydberg atoms
However, the Rutherford-Bohr model had a number of fl aws For example, the spectra of multi-electron atoms had far more lines than the simple Bohr model predicted Nor could the model explain the splitting of the spectral
lines in a magnetic fi eld (a phenomenon known as the Zeeman effect) Within
a short time, a radically different model, the quantum mechanical model, was proposed to account for these observations
n 3
n 2
n 1
FIGURE 1.1 The
Rutherford-Bohr electron-shell model of the
atom, showing the n 5 1, 2, and 3
energy levels.
Aglowing body, such as the Sun, is expected to emit a
continuous spectrum of electromagnetic radiation
However, in the early nineteenth century, a German
sci-entist, Josef von Fraunhofer, noticed that the visible
spec-trum from the Sun actually contained a number of dark
bands Later investigators realized that the bands were
the result of the absorption of particular wavelengths by
cooler atoms in the “atmosphere” above the surface of
the Sun The electrons of these atoms were in the ground
state, and they were absorbing radiation at wavelengths
corresponding to the energies needed to excite them to
higher energy states A study of these “negative” spectra
led to the discovery of helium Such spectral studies are
still of great importance in cosmochemistry—the study
of the chemical composition of stars.
In 1955, two groups of scientists, one in Australia and
the other in Holland, fi nally realized that the absorption
method could be used to detect the presence of elements
at very low concentrations Each element has a lar absorption spectrum corresponding to the various separations of (differences between) the energy levels
particu-in its atoms When light from an atomic emission source
is passed through a vaporized sample of an element, the particular wavelengths corresponding to the various en- ergy separations will be absorbed We fi nd that the higher the concentration of the atoms, the greater the proportion
of the light that will be absorbed This linear relationship between light absorption and concentration is known as
Beer’s law The sensitivity of this method is extremely
high, and concentrations of parts per million are easy to determine; some elements can be detected at the parts per billion level Atomic absorption spectroscopy has now become a routine analytical tool in chemistry, metal- lurgy, geology, medicine, forensic science, and many other
fi elds of science—and it simply requires the movement of electrons from one energy level to another.
Atomic Absorption Spectroscopy
Trang 253
1.1 The Schrödinger Wave Equation and Its Signifi cance
The more sophisticated quantum mechanical model of atomic structure was
derived from the work of Louis de Broglie De Broglie showed that, just as
elec-tromagnetic waves could be treated as streams of particles (photons), moving
particles could exhibit wavelike properties Thus, it was equally valid to picture
electrons either as particles or as waves Using this wave-particle duality, Erwin
Schrödinger developed a partial differential equation to represent the behavior
of an electron around an atomic nucleus One form of this equation, given here
for a one-electron atom, shows the relationship between the wave function of
the electron, C, and E and V, the total and potential energies of the system,
re-spectively The second differential terms relate to the wave function along each
of the Cartesian coordinates x, y, and z, while m is the mass of an electron, and
The derivation of this equation and the method of solving it are in the realm
of physics and physical chemistry, but the solution itself is of great importance
to inorganic chemists We should always keep in mind, however, that the wave
equation is simply a mathematical formula We attach meanings to the solution
simply because most people need concrete images to think about subatomic
phenomena The conceptual models that we create in our macroscopic world
cannot hope to reproduce the subatomic reality
It was contended that the real meaning of the equation could be found
from the square of the wave function, C2, which represents the probability
of fi nding the electron at any point in the region surrounding the nucleus
There are a number of solutions to a wave equation Each solution describes a
different orbital and, hence, a different probability distribution for an
elec-tron in that orbital Each of these orbitals is uniquely defi ned by a set of three
integers: n, l, and m l Like the integers in the Bohr model, these integers are
also called quantum numbers
In addition to the three quantum numbers derived from the original theory,
a fourth quantum number had to be defi ned to explain the results of an
experi-ment in 1922 In this experiexperi-ment, Otto Stern and Walther Gerlach found that
passing a beam of silver atoms through a magnetic fi eld caused about half the
atoms to be defl ected in one direction and the other half in the opposite
direc-tion Other investigators proposed that the observation was the result of two
different electronic spin orientations The atoms possessing an electron with
one spin were defl ected one way, and the atoms whose electron had the
oppo-site spin were defl ected in the oppooppo-site direction This spin quantum number
was assigned the symbol m s
The possible values of the quantum numbers are defi ned as follows:
n, the principal quantum number, can have all positive integer values from
1 to q
1.1 The Schrödinger Wave Equation and Its Signifi cance
Trang 26l, the angular momentum quantum number, can have all integer values from n 2 1 to 0.
m l , the magnetic quantum number, can have all integer values from 1l through 0 to 2l.
m s , the spin quantum number, can have values of 112 and 212.When the value of the principal quantum number is 1, there is only one
possible set of quantum numbers n, l, and m l (1, 0, 0), whereas for a principal quantum number of 2, there are four sets of quantum numbers (2, 0, 0; 2, 1, –1;
2, 1, 0; 2, 1, 11) This situation is shown diagrammatically in Figure 1.2 To identify the electron orbital that corresponds to each set of quantum numbers,
we use the value of the principal quantum number n, followed by a letter for the angular momentum quantum number l Thus, when n 5 1, there is only the 1s orbital.
When n 5 2, there is one 2s orbital and three 2p orbitals (corresponding
to the m l values of 11, 0, and –1) The letters s, p, d, and f are derived from
categories of the spectral lines: sharp, principal, diffuse, and fundamental The correspondences are shown in Table 1.1
When the principal quantum number n 5 3, there are nine sets of quantum numbers (Figure 1.3) These sets correspond to one 3s, three 3p, and fi ve 3d orbitals A similar diagram for the principal quantum number n 5 4 would show 16 sets of quantum numbers, corresponding to one 4s, three 4p, fi ve 4d,
FIGURE 1.2 The possible sets
of quantum numbers for n 5 1
number l and orbital designation
l Value Orbital designation
Trang 271.2 Shapes of the Atomic Orbitals
Representing the solutions to a wave equation on paper is not an easy task In
fact, we would need four-dimensional graph paper (if it existed) to display the
complete solution for each orbital As a realistic alternative, we break the wave
equation into two parts: a radial part and an angular part
Each of the three quantum numbers derived from the wave equation
rep-resents a different aspect of the orbital:
The principal quantum number n indicates the size of the orbital.
The angular momentum quantum number l represents the shape of the
orbital
The magnetic quantum number m l represents the spatial direction of the
orbital
The spin quantum number m s has little physical meaning; it merely allows
two electrons to occupy the same orbital
It is the value of the principal quantum number and, to a lesser extent the
angular momentum quantum number, which determines the energy of the
electron Although the electron may not literally be spinning, it behaves as if it
was, and it has the magnetic properties expected for a spinning particle
An orbital diagram is used to indicate the probability of fi nding an electron
at any point in space We defi ne a location where an electron is most probably
number l and number of orbitals
l Value Number of orbitals
0 1
1 3
2 5
3 7
FIGURE 1.3 The possible sets
of quantum numbers for n 5 3 n
and seven 4f orbitals (Table 1.2) Theoretically, we can go on and on, but as we
will see, the f orbitals represent the limit of orbital types among the elements
of the periodic table for atoms in their electronic ground states
5
Trang 28found as an area of high electron density Conversely, locations with a low
prob-ability are called areas of low electron density
The s Orbitals
The s orbitals are spherically symmetric about the atomic nucleus As the
prin-cipal quantum number increases, the electron tends to be found farther from the nucleus To express this idea in a different way, we say that, as the principal quantum number increases, the orbital becomes more diffuse A unique fea-
ture of electron behavior in an s orbital is that there is a fi nite probability of
fi nding the electron close to, and even within, the nucleus This penetration by
s orbital electrons plays a role in atomic radii (see Chapter 2) and as a means
of studying nuclear structure
Same-scale representations of the shapes (angular functions) of the 1s and 2s orbitals of an atom are compared in Figure 1.4 The volume of a 2s orbital is about four times greater than that of a 1s orbital In both cases, the tiny nucleus
is located at the center of the spheres These spheres represent the region in which there is a 99 percent probability of fi nding an electron The total prob-ability cannot be represented, for the probability of fi nding an electron drops
to zero only at an infi nite distance from the nucleus
The probability of fi nding the electron within an orbital will always be tive (since the probability is derived from the square of the wave function and squaring a negative makes a positive) However, when we discuss the bonding
posi-of atoms, we fi nd that the sign related to the original wave function has tance For this reason, it is conventional to superimpose the sign of the wave
impor-function on the representation of each atomic orbital For an s orbital, the sign
is positive
In addition to the considerable difference in size between the 1s and the 2s orbitals, the 2s orbital has, at a certain distance from the nucleus, a spherical
surface on which the electron density is zero A surface on which the
probabil-ity of fi nding an electron is zero is called a nodal surface When the principal
quantum number increases by 1, the number of nodal surfaces also increases
by 1 We can visualize nodal surfaces more clearly by plotting a graph of the dial density distribution function as a function of distance from the nucleus for
ra-any direction Figure 1.5 shows plots for the 1s, 2s, and 3s orbitals These plots
show that the electron tends to be farther from the nucleus as the principal quantum number increases The areas under all three curves are the same
FIGURE 1.4 Representations of
the shapes and comparative sizes
of the 1s and 2s orbitals.
Trang 297
Electrons in an s orbital are different from those in p, d, or f orbitals in two
signifi cant ways First, only the s orbital has an electron density that varies in
the same way in every direction out from the atomic nucleus Second, there is
a fi nite probability that an electron in an s orbital is at the nucleus of the atom
Every other orbital has a node at the nucleus
The p Orbitals
Unlike the s orbitals, the p orbitals consist of two separate volumes of space
(lobes), with the nucleus located between the two lobes Because there are
three p orbitals, we assign each orbital a direction according to Cartesian
co-ordinates: we have p x , p y , and p z Figure 1.6 shows representations of the three
2p orbitals At right angles to the axis of higher probability, there is a nodal
plane through the nucleus For example, the 2p z orbital has a nodal surface in
the xy plane In terms of wave function sign, one lobe is positive and the other
negative
FIGURE 1.5 The variation of the radial density distribution function with distance from
the nucleus for electrons in the 1s, 2s, and 3s orbitals of a hydrogen atom.
2s
Distance (nm)
0.2 0.4 0.6 0.8 0.2 0.4 0.6 0.8 1.0 1.2Distance (nm)
Trang 30If we compare graphs of electron density as a function of atomic radius for
the 2s orbital and a 2p orbital (the latter plotted along the axis of higher ability), we fi nd that the 2s orbital has a much greater electron density close
prob-to the nucleus than does the 2p orbital (Figure 1.7) Conversely, the second maximum of the 2s orbital is farther out than the single maximum of the 2p
orbital However, the mean distance of maximum probability is the same for both orbitals
Like the s orbitals, the p orbitals develop additional nodal surfaces within the orbital structure as the principal quantum number increases Thus, a 3p orbital does not look exactly like a 2p orbital since it has an additional nodal
surface However, the detailed differences in orbital shapes for a particular angular momentum quantum number are of little relevance in the context of basic inorganic chemistry
The d Orbitals
The fi ve d orbitals have more complex shapes Three of them are located between
the Cartesian axes, and the other two are oriented along the axes In all cases, the nucleus is located at the intersection of the axes Three orbitals each have four lobes that are located between pairs of axes (Figure 1.8) These orbitals
are identifi ed as d xy , d xz , and d yz The other two d orbitals, d z2 and d x22y2, are shown
in Figure 1.9 The d z2 orbital looks somewhat similar to a p z orbital (see Figure 1.6), except that it has an additional doughnut-shaped ring of high electron density
in the xy plane The d x22y2 orbital is identical to the d xy orbital but has been rotated through 458
FIGURE 1.7 The variation of
the radial density distribution
function with distance from the
nucleus for electrons in the 2s and
2p orbitals of a hydrogen atom.
Trang 319
The f Orbitals
The f orbitals are even more complex than the d orbitals There are seven
f orbitals, four of which have eight lobes The other three look like the d z2
orbital but have two doughnut-shaped rings instead of one These orbitals
are rarely involved in bonding, so we do not need to consider them in any
detail
1.3 The Polyelectronic Atom
In our model of the polyelectronic atom, the electrons are distributed among
the orbitals of the atom according to the Aufbau (building-up) principle This
simple idea proposes that, when the electrons of an atom are all in the ground
state, they occupy the orbitals of lowest energy, thereby minimizing the atom’s
total electronic energy Thus, the confi guration of an atom can be described
simply by adding electrons one by one until the total number required for the
element has been reached
Before starting to construct electron confi gurations, we need to take into
account a second rule: the Pauli exclusion principle According to this rule, no
two electrons in an atom may possess identical sets of the four quantum
num-bers Thus, there can be only one orbital of each three-quantum-number set
per atom and each orbital can hold only two electrons, one with m s5 112 and
the other with m s5 212
Filling the s Orbitals
The simplest confi guration is that of the hydrogen atom According to the
Aufbau principle, the single electron will be located in the 1s orbital This
con-fi guration is the ground state of the hydrogen atom Adding energy would raise
the electron to one of the many higher energy states These confi gurations are
referred to as excited states In the diagram of the ground state of the
hydro-gen atom (Figure 1.10), a half-headed arrow is used to indicate the direction of
electron spin The electron confi guration is written as 1s1, with the superscript
1 indicating the number of electrons in that orbital
1.3 The Polyelectronic Atom
1s
FIGURE 1.10 Electron confi guration of a hydrogen atom.
FIGURE 1.9 Representations of the shapes of the 3d x2 2y2 and 3d z2 orbitals.
Trang 32With a two-electron atom (helium), there is a choice: the second electron
could go in the 1s orbital (Figure 1.11a) or the next higher energy orbital, the 2s orbital (Figure 1.11b) Although it might seem obvious that the second electron would enter the 1s orbital, it is not so simple If the second electron entered the 1s orbital, it would be occupying the same volume of space as the electron al- ready in that orbital The very strong electrostatic repulsions, the pairing energy,
would discourage the occupancy of the same orbital However, by occupying
an orbital with a high probability closer to the nucleus, the second electron will experience a much greater nuclear attraction The nuclear attraction is greater
than the inter-electron repulsion Hence, the actual confi guration will be 1s2,although it must be emphasized that electrons pair up in the same orbital only when pairing is the lower energy option
In the lithium atom the 1s orbital is fi lled by two electrons, and the third electron must be in the next higher energy orbital, the 2s orbital Thus, lithium has the confi guration of 1s22s1 Because the energy separation of an s and its corresponding p orbitals is always greater than the pairing energy in a poly- electronic atom, the electron confi guration of beryllium will be 1s22s2 rather
than 1s22s12p1
Filling the p Orbitals
Boron marks the beginning of the fi lling of the 2p orbitals A boron atom has
an electron confi guration of 1s22s22p1 Because the p orbitals are degenerate
(that is, they all have the same energy), it is impossible to decide which one of the three orbitals contains the electron
Carbon is the second ground-state atom with electrons in the p orbitals
Its electron confi guration provides another challenge There are three
pos-sible arrangements of the two 2p electrons (Figure 1.12): (a) both electrons
in one orbital, (b) two electrons with parallel spins in different orbitals, and (c) two electrons with opposed spins in different orbitals On the basis of elec- tron repulsions, the fi rst possibility (a) can be rejected immediately The deci-
sion between the other two possibilities is less obvious and requires a deeper knowledge of quantum theory In fact, if the two electrons have parallel spins, there is a zero probability of their occupying the same space However, if the spins are opposed, there is a fi nite possibility that the two electrons will occupy the same region in space, thereby resulting in some repulsion and a
higher energy state Hence, the parallel spin situation (b) will have the lowest
energy This preference for unpaired electrons with parallel spins has been
formalized in Hund’s rule: When fi lling a set of degenerate orbitals, the
num-ber of unpaired electrons will be maximized and these electrons will have parallel spins
After the completion of the 2p electron set at neon (1s22s22p6), the 3s and 3p
orbitals start to fi ll Rather than write the full electron confi gurations, a ened form can be used In this notation, the inner electrons are represented by the noble gas symbol having that confi guration Thus, magnesium, whose full
short-electron confi guration would be written as 1s22s22p63s2, can be represented as
having a neon noble gas core, and its confi guration is written as [Ne]3s2 An
FIGURE 1.11 Two possible
electron confi gurations for helium.
2s 2s
1s 1s
Trang 3311
advantage of the noble gas core representation is that it emphasizes the
outer-most (valence) electrons, and it is these electrons that are involved in chemical
bonding Then fi lling the 3p orbitals brings us to argon.
Filling the d Orbitals
It is at this point that the 3d and 4s orbitals start to fi ll The simple orbital
energy level concept breaks down because the energy levels of the 4s and 3d
orbitals are very close What becomes most important is not the minimum
energy for a single electron but the confi guration that results in the least
number of inter-electron repulsions for all the electrons For potassium, this is
[Ar]4s1; for calcium, [Ar]4s2
In general, the lowest overall energy for each transition metal is obtained
by fi lling the s orbitals fi rst; the remaining electrons then occupy the d orbitals
Although there are minor fl uctuations in confi gurations throughout the d-block
and f-block elements, the order of fi lling is quite consistent:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
Figure 1.13 shows the elements organized by order of orbital fi lling
This order is shown as an energy-level diagram in Figure 1.14 (page 13)
The orbitals fi ll in this order because the energy differences between the s, p, d,
and f orbitals of the same principal quantum number become so great beyond
n 5 2 that they overlap with the orbitals of the following principal quantum
numbers It is important to note that Figure 1.14 shows the fi lling order, not
the order for any particular element For example, for elements beyond zinc,
electrons in the 3d orbitals are far lower in energy than those in the 4s orbitals
Thus, at this point, the 3d orbitals have become “inner” orbitals and have no
role in chemical bonding Hence, their precise ordering is unimportant
Although these are the generalized rules, to illustrate how this delicate
balance changes with increasing numbers of protons and electrons, the outer
electrons in each of the Group 3 to Group 12 elements are listed here These
confi gurations are not important in themselves, but they do show how close the
ns and (n – 1)d electrons are in energy.
1.3 The Polyelectronic Atom
Atom Confi guration Atom Confi guration Atom Confi guration
Trang 34s-Block
LiNaCaKRbCs
Fr RaBaSr
MgBe
H He
KrBrSeAsGeGaAl
NeFONCB
ArClSPSi
Uuo Uuh
Uup Uuq Uut
XeAtPo
FIGURE 1.13 In this version of the periodic table, the chemical elements are organized in order of orbital fi lling.
Trang 35For certain elements, the lowest energy is obtained by shifting one or both
of the s electrons to d orbitals Looking at the fi rst series in isolation would
lead to the conclusion that there is some preference for a half-full or full set of
d orbitals by chromium and copper However, it is more accurate to say that
the inter-electron repulsion between the two s electrons is suffi cient in several
cases to result in an s1 confi guration
For the elements from lanthanum (La) to ytterbium (Yb), the situation is
even more fl uid because the 6s, 5d, and 4f orbitals all have similar energies For
example, lanthanum has a confi guration of [Xe]6s25d1, whereas the next
ele-ment, cerium, has a confi guration of [Xe]6s24f2 The most interesting electron
confi guration in this row is that of gadolinium, [Xe]6s25d14f7, rather than the
predicted [Xe]6s24f8 This confi guration provides more evidence of the
impor-tance of inter-electron repulsion in the determination of electron confi guration
1.3 The Polyelectronic Atom
Trang 36when adjacent orbitals have similar energies Similar complexities occur among
the elements from actinium (Ac) to nobelium (No), in which the 7s, 6d, and 5f
orbitals have similar energies
1.4 Ion Electron Confi gurations
For the early main group elements, the common ion electron confi gurations can be predicted quite readily Thus, metals tend to lose all the electrons in the
outer orbital set This situation is illustrated for the isoelectronic series (same
electron confi guration) of sodium, magnesium, and aluminum cations:
Atom Electron confi guration Ion Electron confi guration
by the loss of the 6p electrons only (the “inert pair” effect, which we discuss
in Chapter 9, Section 9.8), whereas the 41 charge results from loss of both 6s and 6p electrons:
Atom Electron confi guration Ion Electron confi guration
Pb [Xe]6s24f145d106p2 Pb21 [Xe]6s24f145d10
Notice that the electrons of the higher principal quantum number are lost fi rst
This rule is found to be true for all the elements For the transition metals, the s
electrons are always lost fi rst when a metal cation is formed In other words, for
the transition metal cations, the 3d orbitals are always lower in energy than the 4s orbitals, and a charge of 21, representing the loss of the two s electrons, is
common for the transition metals and the Group 12 metals For example, zinc always forms an ion of 21 charge:
Atom Electron confi guration Ion Electron confi guration
Zn [Ar]4s23d10 Zn21 [Ar]3d10
Trang 3715
Iron forms ions with charges of 21 and 31 and, as shown here, it is tempting to
ascribe the formation of the 31 ion to a process in which inter-electron
repul-sion “forces out” the only paired d electron:
1.5 Magnetic Properties of Atoms
Atom Electron confi guration Ion Electron confi guration
It is dangerous, however, to read too much into the electron confi gurations of
atoms as a means of predicting the ion charges The series of nickel, palladium,
and platinum illustrate this point: they have different confi gurations as atoms,
yet their common ionic charges and corresponding ion electron confi gurations
are similar:
Atom Electron confi guration Ion Electron confi guration
Ni [Ar]4s23d8 Ni21 [Ar]3d8
Pd [Kr]5s04d10 Pd21, Pd41 [Kr]4d8, [Kr]4d6
Pt [Xe]6s15d9 Pt21, Pt41 [Xe]5d8, [Xe]5d6
1.5 Magnetic Properties of Atoms
In the discussions of electron confi guration, we saw that some atoms possess
unpaired electrons The presence of unpaired electrons in the atoms of an
ele-ment can be determined easily from the eleele-ment’s magnetic properties If atoms
containing only spin-paired electrons are placed in a magnetic fi eld, they are
weakly repelled by the fi eld This phenomenon is called diamagnetism
Con-versely, atoms containing one or more unpaired electrons are attracted by the
magnetic fi eld This behavior of unpaired electrons is named paramagnetism.
The attraction of each unpaired electron is many times stronger than the
repul-sion of all the spin-paired electrons in that atom
To explain paramagnetism in simple terms, we can visualize the electron
as a particle spinning on its axis and generating a magnetic moment, just as
an electric current fl owing through a wire does This permanent magnetic
mo-ment results in an attraction into the stronger part of the fi eld When electrons
have their spins paired, the magnetic moments cancel each other As a result,
the paired electrons are weakly repelled by the lines of force of the magnetic
fi eld In paramagnetic materials, application of a magnetic fi eld aligns some
of the normally randomly oriented electron spins with the applied magnetic
fi eld (Figure 1.15a and b) It is this alignment that results in the attraction of
the material into the magnetic fi eld We will encounter this phenomenon again
in our discussions of covalent bonding and the bonding in transition metal
compounds
Trang 381.6 Medicinal Inorganic Chemistry: An Introduction
Inorganic chemistry affects our lives directly in two ways First, as we discuss
at the end of later chapters, many chemical elements are required for the tioning of living organisms Second, inorganic elements and compounds have been used as medicines since earliest times Periodically in this text, we give examples of the use of inorganic compounds as medicinal substances, but it is useful to provide an overview
func-Many inorganic compounds have been used as medicines through the ages
A fashionable habit in European countries was to “drink the waters” at spa cities In some cases, the springs were mineral-rich; for example, the water in Vichy, France (now available bottled), is rich in magnesium ion, which acts as a potent laxative That water, therefore, should only be drunk in small quantities The solid salt, magnesium sulfate heptahydrate, MgSO4.7H2O, has the same effect It was named Epsom salts after the town in England where it was fi rst discovered During the nineteenth century, one British hospital was using 2.5 tonnes per year on its patients!
Some cultures practice geophagy, the eating of soil—usually clay Clays are a complex class of minerals, as we discuss in Chapter 14 One form of clay
is kaolin—a substance that is known for its absorptive abilities Several types
of tablets to combat stomach upsets employ kaolin, which, it is believed, can surface-absorb toxins produced by ingested harmful bacteria Other clays and soils can supply trace nutrients However, persistent clay eating is not advised since the clay can line the stomach and prevent nutrient adsorption Also, many natural clays contain high concentrations of harmful elements, such as lead
Inorganic medicinal chemistry can appear in the most unusual contexts For example, religious statues made from the mineral realgar, diarsenic disulfi de (As2S2), were popular among devotees of the Chinese Taoist religion Han-dling the statues was believed to restore health In this particular case, chemis-try rather than faith might have contributed, for many people in tropical areas suffer from internal parasites and handling the statues would result in arsenic absorption through the skin, enough to kill the parasites but not enough to kill the devotee
FIGURE 1.15 The behavior of
paramagnetic materials without
(a) and with (b) an applied
magnetic fi eld.
(a)
(b)
H
Trang 3917
In the following chapters, we mention a few of the many modern medicinal
applications of inorganic compounds:
Antacids (Chapter 7)
Lithium in the treatment of bipolar disorder (Chapter 9)
Boron neutron capture therapy (Chapter 13)
Platinum complexes as anticancer agents (Chapter 19)
Technetium as a radiopharmaceutical (Chapter 21)
Gold in the treatment of rheumatoid arthritis (Chapter 21)
KEY IDEAS
• The properties of an electron in an atom can be defi ned
in terms of four quantum numbers
• There are a variety of shapes of orbitals (s, p, d, f ) as
defi ned by the angular momentum quantum number.
• Electrons in the ground state of an atom fi ll the orbitals
of lowest energy.
• For the transition metals, the energies of the ns and (n – 1)d orbitals are very similar.
• In the formation of a cation, the electrons in the orbitals
of highest principal quantum number are lost fi rst.
• Paramagnetic behavior in a magnetic fi eld indicates the presence of unpaired electrons.
EXERCISES
1.1 Defi ne the following terms: (a) nodal surface;
(b) Pauli exclusion principle; (c) paramagnetic.
1.2 Defi ne the following terms: (a) orbital; (b) degenerate;
(c) Hund’s rule.
1.3 Construct a quantum number tree for the principal
quantum number n 5 4 similar to that depicted for n 5 3 in
Figure 1.3.
1.4 Determine the lowest value of n for which ml can
(theoretically) have a value of 14.
1.5 Identify the orbital that has n 5 5 and l 5 1.
1.6 Identify the orbital that has n 5 6 and l 5 0.
1.7 How does the quantum number n relate to the
properties of an orbital?
1.8 How does the quantum number l relate to the
properties of an orbital?
1.9 Explain concisely why carbon has two electrons in
different p orbitals with parallel spins rather than the other
possible arrangements.
1.10 Explain concisely why beryllium has a ground-state
electron confi guration of 1s22s2 rather than 1s22s12p1.
1.11 Write noble gas core ground-state electron confi urations for atoms of (a) sodium; (b) nickel; (c) copper.
g-1.12 Write noble gas core ground-state electron confi urations for atoms of (a) calcium; (b) chromium; (c) lead.
g-1.13 Write noble gas core ground-state electron confi rations for ions of (a) potassium; (b) scandium 31; (c) cop- per 21.
gu-1.14 Write noble gas core ground-state electron confi rations for ions of (a) chlorine; (b) cobalt 21; (c) manga- nese 41.
gu-1.15 Predict the common charges of the ions of thallium Explain your reasoning in terms of electron confi gurations.
1.16 Predict the common charges of the ions of tin Explain your reasoning in terms of electron confi gurations.
1.17 Predict the common charge of the silver ion Explain your reasoning in terms of electron confi gurations.
1.18 Predict the highest possible charge of a zirconium ion Explain your reasoning in terms of electron confi gurations.
1.19 Use diagrams similar to Figure 1.12 to determine the number of unpaired electrons in atoms of (a) oxygen; (b) magnesium; (c) chromium.
Exercises
Trang 40BEYOND THE BASICS
1.23 The next set of orbitals after the f orbitals are the g
orbitals How many g orbitals would there be? What would
be the lowest principal quantum number n that would
pos-sess g orbitals? Deduce the atomic number of the fi rst
ele-ment at which g orbitals would begin to be fi lled on the
basis of the patterns of the d and f orbitals.
1.24 An alternative to the Schrödinger wave equation is
the Dirac wave equation Using online sources, research the
Dirac wave equation and contrast it with the Schrödinger
wave equation.
1.25 Use an advanced inorganic chemistry text as a
source of information on the f orbitals What are their
com-mon features? How do they differ acom-mong themselves?
1.26 In Section 1.3, gadolinium is mentioned as having
an electron confi guration that deviates from the lanthanoid pattern Which element in the actinoids should show a simi- lar deviation? What would be its electron confi guration?
1.27 In Figure 1.13, the elements are organized logically according to the order of orbital fi lling Identify two disad- vantages of organizing the elements in this way.
1.28 A philosophical question: Does an orbital exist even
if it does not contain an electron? Discuss.
1.20 Use diagrams similar to Figure 1.12 to determine
the number of unpaired electrons in atoms of (a) nitrogen;
(b) silicon; (c) iron.
1.21 Write the electron confi guration expected for
ele-ment 113 and the confi gurations for the two cations that it
is most likely to form.
1.22 Which of the following species are hydrogen-like? (a) He 1 ; (b) He – ; (c) Li 1 ; (d) Li 21
ADDITIONAL RESOURCES
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