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The easy way to get a grip on inorganic chemistry
Inorganic chemistry can be an intimidating subject, but it doesn't have to be! Whether you're currently enrolled in an inorganic chemistry class or you have a background in chemistry and want to expand your knowledge, Inorganic Chemistry For Dummies is the approachable, hands-on guide you can trust for fast, easy learning.
Inorganic Chemistry For Dummies features a thorough introduction to the study of the synthesis and behavior of inorganic and organometallic compounds. In plain English, it explains the principles of inorganic chemistry and includes worked-out problems to enhance your understanding of the key theories and concepts of the field.
If you're pursuing a career as a nurse, doctor, or engineer or a lifelong learner looking to make sense of this fascinating subject, Inorganic Chemistry For Dummies is the quick and painless way to master inorganic chemistry.
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Seitenzahl: 575
Veröffentlichungsjahr: 2013
Inorganic Chemistry For Dummies®
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ISBN 978-1-118-21794-8 (pbk); ISBN 978-1-118-22882-1 (ebk); ISBN 978-1-118-22891-3 (ebk); ISBN 978-1-118-22894-4 (ebk)
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About the Authors
Michael L. Matson started studying chemistry at the U.S. Naval Academy in Annapolis, Maryland. After leaving the Navy, Michael started a PhD program at Rice University, studying the use of carbon nanotubes for medical diagnosis and treatment of cancer. Specifically, Michael focused on internalizing radioactive metal ions within carbon nanotubes: Some radioactive metals could be pictured with special cameras for diagnosis, whereas others were so powerful they could kill cells for treatment. It was at Rice that Michael and Alvin met. Following Rice, Michael went to the University of Houston-Downtown to begin a tenure-track professorship. Happily married to a woman he first met in seventh grade, Michael has two young children, a yellow Labrador retriever named Flounder, is a volunteer firefighter and sommelier, and enjoys CrossFitting.
Alvin W. Orbaek was introduced to chemistry at Rice University (Houston, Texas) by way of nanotechnology, where he studied single-walled carbon nanotubes, transition metal catalysts, and silver nanoparticles. He had previously received a degree in Experimental Physics from N.U.I. Galway (Ireland) and moved into the study of space science and technology at the International Space University (Strasbourg, France). He received a position on Galactic Suite, an orbiting space hotel. To date, he enjoys life by sailing, snowboarding, and DJing. He has been spinning vinyl records since the Atlantic Hotel used to rave, and the sun would set in Ibiza. He hopes to empower people through education and technology, to that effect he is currently completing a PhD in Chemistry at Rice University.
Dedications
Michael: To my wife, Samantha.
Alvin: To Declan, Ann Gitte, Anton, Anna-livia, and Bedstemor.
Authors’ Acknowledgments
Michael: I’d like to acknowledge the immeasurable amounts of assistance from Matt Wagner, Susan Hobbs, Lindsay Lefevere, Alecia Spooner, and Joan Freedman.
Alvin: Without John Wiley & Sons, there would be no book, and for that I am very grateful. Particularly because of the very positive and professional attitude by which they carry out their business; thanks for getting it done. It was a blessing to work with you. In particular, I would like to mention Alecia Spooner, Susan Hobbs (Suz), and Lindsay Lefevere, and thanks to the technical editors (Reynaldo Barreto and Bradley Fahlman) for their crucial input. I would also like to thank Matt Wagner for invaluable support and assistance. And to Mike Matson, thank you for the invitation to write this book.
I have had many teachers, mentors, and advisors throughout the years, but there are five who deserve attention. Andrew Smith at Coleenbridge Steiner school, where I enjoyed learning a great deal. John Treacy, who made every science class the most riveting class each day. Pat Sweeney, whose habit of teaching would leave anyone engrossed in mathematics. To Ignasi Casanova for his mentorship and introduction to the nanos. And Andrew Barron, both my PhD advisor and mentor, to whom I owe a great deal of credit, due in no small part to his measure of tutelage.
But all this stands upon a firm foundation that is based on the support of Dec, Gitte, Anton, and Anna; here’s to next Christmas — whenever. There are many other friends and family who have contributed to this work, too many to mention them all. But I’d especially like to thank my colleagues from the Irish house, who so graciously agreed to read through the text, namely Alan Taylor, Nigel Alley, and Stuart Corr. Also to Sophia Phounsavath and Brandon Cisneros for proofreading. Jorge Fallas for the Schrödinger equation. To Gordon Tomas for continued support of my writing. And to Gabrielle Novello, who fed me wholesome foods while I otherwise converted coffee and sleepless nights into this book. And to Valhalla for those nights when work was not working for me. And to PHlert, the best sailing program on this planet, or any other.
Publisher’s Acknowledgments
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Table of Contents
Introduction
About This Book
Conventions Used in This Book
What You Don’t Need to Read
Foolish Assumptions
How This Book Is Organized
Part I: Reviewing Some General Chemistry
Part II: Rules of Attraction: Chemical Bonding
Part III: It’s Elemental: Dining at the Periodic Table
Part IV: Special Topics
Part V: The Part of Tens
Icons Used in This Book
Where to Go from Here
Part I: Reviewing Some General Chemistry
Chapter 1: Introducing Inorganic Chemistry
Building the Foundation
Losing your electrons
Splitting atoms: Nuclear chemistry
Changing pH
Getting a Grip on Chemical Bonding
Traveling Across the Periodic Table
Hyping up hydrogen
Moving through the main groups
Transitioning from one side of the table to another
Uncovering lanthanides and actinides
Diving Deeper: Special Topics
Bonding with carbon: Organometallics
Speeding things up: Catalysts
Inside and out: Bio-inorganic and environmental chemistry
Solid-state chemistry
Nanotechnology
Listing 40 More
Chapter 2: Following the Leader: Atomic Structure and Periodic Trends
Up an’ Atom: Reviewing Atomic Terminology
Sizing up subatomic particles
Knowing the nucleus
Going orbital
Distinguishing atomic number and mass number
Identifying isotopes
Grouping Elements in the Periodic Table
Keeping up with periodic trends
Measuring atomic size
Rating the atomic radius
Eyeing ionization energy
Examining electron affinities
Noting electronegativity
Chapter 3: The United States of Oxidation
Entering the Oxidation-Reduction Zone
Following oxidation state rules
Scouting reduction potentials
Walking through a Redox Reaction
Isolating Elements
Mechanically separating elements
Using thermal decomposition
Displacing one element with another
Heating things up: High-temperature chemical reactions
Relying on electrolytic reduction
Chapter 4: Gone Fission: Nuclear Chemistry
Noting Nuclear Properties
Using the force
The empirical strikes back
Documenting Atomic Decay: Radioactivity
Alpha radiation
Beta radiation
Gamma radiation
The half-life principle
Blind (radiocarbon) dating
Radioisotopes
Catalyzing a Nuclear Reaction
Fission
Fusion
Chapter 5: The ABCs: Acid-Base Chemistry
Starting with the Basics: Acids and Bases
Developing the pH Scale
Calculating pH
Calculating acid dissociation
Touring Key Theories: A Historical Perspective
The early years
Brønsted-Lowry theory
Accepting or donating: Lewis’s theory
Comparing Lewis and Brønsted theories
Pearson’s Hard and Soft Acids and Bases (HSAB)
Characterization of the hard bodies
Who you callin’ soft?
Strapping on a Cape: Superacids
Part II: Rules of Attraction: Chemical Bonding
Chapter 6: No Mr. Bond, I Expect You to π: Covalent Bonding
Connecting the Dots: Lewis Structures
Counting electrons
Placing electrons
Price tags in black ties? Formal charges
Returning to the drawing board: Resonance structures
Keeping Your Distance: VSEPR
Ante Up One Electron: Valence-Bond Theory
Summing It All Up: Molecular Orbital Theory
Types of MOs
Evens and odds: Gerade and ungerade symmetry
Identical twins: Homonuclear diatomic molecules
Fraternal twins: Heteronuclear diatomic molecules
Chapter 7: Molecular Symmetry and Group Theory
Identifying Molecules: Symmetry Elements and Operations
Identity
n-fold rotational axis
Inversion center
Mirror planes
Improper rotation axis
It’s Not Polite to Point! Molecular Point Groups
Being Such a Character Table
Dissecting a character table
Degrees of freedom
A glitch in the matrix: Matrix math
Reducible reps
Infrared and Raman active modes
Chapter 8: Ionic and Metallic Bonding
Blame It on Electrostatic Attraction: Forming Ionic Bonds
Marrying a cation and an anion
Measuring bond strength: Lattice energy
Coexisting with covalent bonds
Conducting electricity in solution
Admiring Ionic Crystals
Studying shapes: Lattice types
Size matters (when it’s ionic)
“I’m Melting!” Dissolving Ionic Compounds with Water: Solubility
Just add water: Hydrated ions
Counting soluble compounds
What Is a Metal, Anyway?
Tracing the history of metallurgy
Admiring the properties of solid metals
Delocalizing electrons: Conductivity
Analyzing alloys
Swimming in the Electron Sea: Metallic Bonding Theories
Free-electron theory
Valence bond theory
Band theory
Chapter 9: Clinging to Complex Ions: Coordination Complexes
Counting bonds
Seeking stability
Grouping geometries
Identifying Isomers
Connecting differently: Structural isomers
Arranged differently: Stereoisomers
Naming Coordination Complexes
Sorting Out the Salts
Creating Metal Complexes throughout the Periodic Table
Alkali metals
Alkali earth metals
Transition metals
Lanthanides and actinides
Metalloids
Applying Coordination Complexes in the Real World
Part III: It’s Elemental: Dining at the Periodic Table
Chapter 10: What the H? Hydrogen!
Visiting Hydrogen at Home: Its Place in the Periodic Table
Appreciating the Merits of Hydrogen
Available in abundance
Molecular properties
Nuclear spin
Introducing Hydrogen Isotopes
Investing in Hydrogen Bonds
Forming a hydrogen ion
Creating hydrides
Applying Itself: Hydrogen’s Uses in Chemistry and Industry
Chapter 11: Earning Your Salt: The Alkali and Alkaline Earth Metals
Salting the Earth: Group 1 Elements
Lithium the outlier
Seafaring sodium
Maintaining your brain with potassium
Rubidium, cesium, francium, oh my
Reacting Less Violently: The Group 2 Alkaline Earth Metals
Being beryllium
Magnificent magnesium
Commonly calcium
Strontium, barium, radium
Diagramming the Diagonal Relationship
Chapter 12: The Main Groups
Placing Main Group Elements on the Periodic Table
Lucky 13: The Boron Group
Not-so-boring boron
An abundance of aluminum
Mendeleev’s Missing Link: Gallium
Increasing indium use
Toxic thallium
The Diamond Club: The Carbon Group
Captivating carbon
Coming in second: Silicon
Germane germanium
Malleable tin cans
Plumbing lead
Noting Pnictides of the Nitrogen Group
Leading the pnictides: Nitrogen
Finding phosphorus everywhere
Melding the metalloids: Arsenic and antimony
Keeping Up with the Chalcogens
Oxygen all around
Sulfur
From the Earth to the moon
Marco — polonium!
(Re)Active Singles: The Group 17 Halogens
Cleaning up with chlorine
Briny bromine
Iodine
Rarely astatine
Lights of New York: The Group 18 Noble Gases
Chapter 13: Bridging Two Sides of the Periodic Table: The Transition Metals
Getting to Know Transition Metals
Sorting T-metals into series
Separating T-metals from the main group
Partially Filling d-Orbitals
Calculating an effective nuclear charge
Forming more than one oxidation state
Splitting the Difference: Crystal Field Theory and Transition Metal Complexes
Dividing d-orbitals
Absorbing light waves: Color
Building attraction: Magnetism
Electronic Structure and Bonding
Reacting with other elements
Creating coordination complexes
Adsorbing gas: T-metals in catalysis
Chapter 14: Finding What Lies Beneath: The Lanthanides and Actinides
Spending Quality Time with the Rare Earth Elements: Lanthanides
Electronic structure
Reactivity
Lanthanide contraction
Separating the lanthanide elements
Using lanthanides
Feelin’ Radioactive: The Actinides
Finding or making actinides
Examining electronic structure
Comparing Reactivity: Actinide versus Lanthanide
Looking More Closely at Uranium
Part IV: Special Topics
Chapter 15: Not Quite Organic, Not Quite Inorganic: Organometallics
Building Organometallic Complexes
Adhering to Electron Rules
Counting to eight: The octet rule
Calculating with the 18-electron rule
Settling for 16 electrons
Effectively using the EAN rule
Bonding with Metals: Ligands
Including Carbon: Carbonyls
Providing the Best Examples
e-precise carbon
e-rich nitrogen
e- deficient boron
Behaving Oddly: Organometallics of Groups 1, 2, and 12
Sandwiched Together: Metallocenes
Clustering Together: Metal-Metal Bonding
Creating Vacancies: Insertion and Elimination
Synthesizing Organometallics
Showing Similarities with Main Group Chemistry
Chapter 16: Accelerating Change: Catalysts
Speeding Things Up – The Job of a Catalyst
Considering Types of Catalysts
Homogenous catalysts
Heterogeneous
Organocatalysts
Chapter 17: Bioinorganic Chemistry: Finding Metals in Living Systems
Focusing on Photosynthesis
Climbing Aboard the Oxygen Transport
Feeding a Nitrogen Fixation
Fixing nitrogen for use by organisms
Re-absorbing nitrogen
Being Human
Making things happen: Enzymes
Curing disease: Medicines
Causing problems: Toxicity
Answering When Nature Calls: Environmental Chemistry
Eyeing key indicators
Rocking the heavy metals
Killing me softly: Pesticides
Looking for and removing contaminants
Chapter 18: Living in a Materials World: Solid-State Chemistry
Studying Solid Structures
Building crystals with unit cells
Labeling lines and corners: Miller indices
Three Types of Crystal Structure
Simple crystal structures
Binary crystal structures
Complex crystal structures
Calculating Crystal Formation: The Born-Haber Cycle
Bonding and Other Characteristics
Characterizing size
Dissolving in liquids: Solubility
Encountering zero resistance: Superconductivity
Information technology: Semiconductors
Synthesizing Solid Structures
Detecting Crystal Defects
Chapter 19: Nanotechnology
Defining nanotechnology
History of nanotechnology
The science of nanotechnology
Top-down versus bottom-up
Nanomaterials
Size and shape control
Self-assembly and gray goo
Applications for Nanotechnology
Cancer therapy
Catalysis
Education
Part V: The Part of Tens
Chapter 20: Ten Nobels
Locating Ligands: Alfred Werner
Making Ammonia: Fritz Haber
Creating Transuranium Elements: McMillan and Seaborg
Adding Electronegativity: Pauling
Preparing Plastics: Ziegler and Natta
Sandwiching Compounds: Fischer and Wilkinson
Illuminating Boron Bonds: Lipscomb
Characterizing Crystal Structures: Hauptman and Karle
Creating Cryptands: Jean-Marie Lehn
Making Buckyballs
Chapter 21: Tools of the Trade: Ten Instrumental Techniques
Absorbing and Transmitting Light Waves: UV-vis and IR
Catching Diffracted Light: XRD
Rearranging Excited Atoms: XRF
Measuring Atoms in Solution: ICP/AA
Detecting Secondary Electrons: SEM
Reading the Criss-Crossed Lines: TEM
Characterizing Surface Chemistry: XPS
Evaporating Materials: TGA
Cyclic Voltammetry
Tracking Electron Spin: EPR
Chapter 22: Ten Experiments
Turning Blue: The Clock Reaction
Forming Carbon Dioxide
The Presence of Carbon Dioxide
Mimicking Solubility
Separating Water into Gas
Testing Conductivity of Electrolyte Solutions
Lemon Batteries
Purifying Hydrogen
Colorful Flames
Making Gunpowder
Chapter 23: Ten Inorganic Household Products
Salting Your Food
Bubbling with Hydrogen Peroxide
Baking with Bicarbonate
Whitening with Bleach
Using Ammonia in Many Ways
Killing Pests with Borax
Soothing Babies with Talc
Cleaning with Lye
Scratching Stainless Steel
Wrapping It Up with Aluminum Foil
Glossary
Introduction
Inorganic chemistry deals with all the atoms on the periodic table, the various rules that govern how they look, and how they interact. At first glance, trying to understand the differences among 112 atoms might seem like a mammoth task. But because of the periodic table, we can bunch them up into groups and periods and make them much easier to grasp.
So welcome to Inorganic Chemistry For Dummies. We hope that through this book you come to learn a great deal about the environment around you, what materials you use on a regular basis, and why some materials are more important to us than others. This book is fun and informative, while at the same time insightful and descriptive. And it’s designed to make this fascinating and practical science accessible to anyone, from the novice chemist to the mad scientist.
About This Book
This book was written in such a way that you can start in any chapter you choose, in the chapter that interests you the most, without having to read all the chapters before it. But the chapters build on material from one chapter to the next, so if you feel more background would help you, feel free to start with Chapter 1. You can also make use of the numerous cross references in each chapter to find pertinent information. But it can also be read like a study guide to help a student understand some of the more complicated aspect of this fascinating science.
We tried to make the information as accessible as possible. Each chapter is broken down into bite-sized chunks that make it easy for you to quickly digest and understand the material presented. Some of the chunks are further broken down into subsections when there’s special need to elaborate further on the concepts being discussed.
Science is a process that requires lots of imagination. It requires more imagination than memory, especially as you start to learn more and more about a certain topic. To help with your imagination we have tried to include helpful graphics and artwork that complement the writing within the text. Further to this we include many real-world examples and interesting historical or scientific tidbits to keep your curiosity piqued.
Conventions Used in This Book
Science progressed more rapidly in the last 200 years than it had in the few thousand years previous. A great deal of this success came from the agreement among scientist to create and use a set of standard conventions. The two most important conventions are the periodic table and the international system of units, called SI units. SI units are based on the metric system, and it’s more common to see temperature expressed as Celsius than Fahrenheit. And you see lengths expressed in meters instead of inches and feet. Weights and mass are expressed in terms of grams instead of pounds or stone.
And the following conventions throughout this text make everything consistent and easy to understand:
All Web addresses appear in monofont.
New and key terms appear in italics and are closely followed by an easy-to-understand definition.
Bold text highlights the action part of numbered steps.
What You Don’t Need to Read
Sidebars are highlighted in gray-shaded boxes so they’re easy to pick out. They contain fun facts and curious asides, but none of their information is crucial to your understanding of inorganic chemistry. Feel free to just skip over them if you prefer.
Foolish Assumptions
As authors of Inorganic Chemistry For Dummies we may have made a few foolish assumptions about the readership. We assume that you have very little background in chemistry, and possibly none at all; that you’re new to inorganic chemistry, and maybe you have never heard of the subject before. We assume that you know what chemistry is, but not much more than that. This book begins with all the general chemistry info that you need to grasp the concepts and material in the rest of the book. If you have some understanding of general chemistry, however, all the better.
You may be a medical student who needs to brush up in inorganic chemistry, or a high school student getting ready for a science fair, or even a freshman or junior at college. We’ve tailored this book to meet all your needs, and we sincerely hope you find great explanations about the concepts presented that are also engaging, interesting, and useful.
When you finish reading this book and your interest in chemistry is heightened, we recommend that you go to a local bookseller (second-hand book stores are a personal favorite) and find more books that offer other perspectives on inorganic chemistry. There are also excellent resources on the Internet, and many schools make class notes available online. But the best way to get involved in chemistry is by doing it. Chemistry is a fun and exciting field, made evident when you conduct chemistry experiments. Keep an eye out for demonstration kits that enable you to do your own experiments at home. And note that the last chapter of this book offers ten really cool experiments, too.
How This Book Is Organized
This book is organized into multiple parts that group topics together in the most logical way possible. Here’s a brief description of each section of Inorganic Chemistry For Dummies:
Part I: Reviewing Some General Chemistry
Here you are introduced to science in general, and we give you the basic tenets of general chemistry that help you throughout the rest of the book.
In Chapter 1, you start with an introduction to inorganic chemistry, what it is, and why it is important. You learn how it’s different from organic chemistry and how this difference is important for technology and society.
The following chapters of this section deal with topics that are covered in many general chemistry textbooks, but these chapters cover the topics in greater detail than a general chemistry textbook. In Chapter 2 we explain what the atom looks like, how it’s structured, and why this is important for inorganic chemistry. In particular, this chapter delves into the periodic table and how the structure of the atom is described. Chapter 3 introduces oxidation and reduction chemistry that helps you understand why many chemical reactions take place. It deals with the electrons that each atom has and how the electrons can be shuttled around from atom to atom. Then in Chapter 4 we focus on the nucleus and how changes to the nucleus lead to nuclear chemistry. And finally we end this section by talking about acid-base chemistry because this can help you understand the many ways in which atoms and molecules interact with one another.
Part II: Rules of Attraction: Chemical Bonding
In this section we talk about the various ways that atoms can bond with one another. In Chapter 6 we introduce covalent bonding. Chapter 7 deals with molecular symmetry, not just for inorganic chemistry but also fundamental to many of the physical sciences. Ionic and metallic bonding are detailed in Chapter 8.
Chapter 9, like all of the chapters, can be read as a standalone chapter, but it’s much easier to understand if you read through the three preceding chapters. If you get stuck on coordination complexes, however, refer back to the previous three chapters for a little background information.
Part III: It’s Elemental: Dining at the Periodic Table
The periodic table contains over 100 separate and unique elements, which are described in Part III. We cover all the important elements; and to make it easier to digest, we’ve broken them down into five related chapters. Each chapters deals with elements that are similar to each other, making them easier to understand.
To get the ball rolling we introduce hydrogen in Chapter 10, because it’s the most abundant element in the universe and can be found in many chemicals and materials. We then move from left to right on the periodic table, starting off with the alkali and alkali earth elements in Chapter 11. We guide you through the periodic table to the main group elements in Chapter 12, the transition metals in Chapter 13, and finally round out Part III with the lanthanides and actinides in Chapter 14.
Part IV: Special Topics
These chapters cover what makes the study of inorganic chemistry so interesting and also distinguishes it from organic chemistry. However, you will find a great deal of overlap with other fields of study such as material science, physics, and biology.
Inorganic chemistry became a modern science with the advent of organometallic chemistry, described in Chapter 15. Chapter 16 shows you how practical and important catalysis is to the modern world in which we live. Chapter 17 deals with the inorganic chemistry of living systems and the environment. The subject matter makes this chapter unique from the others in this section. This is also true for Chapter 18 where we describe solid state chemistry, the basis of the information technology revolution. Chapter 19 gives you a quick introduction to one of the most interesting and promising technological developments of the modern age, namely nanotechnology.
Part V: The Part of Tens
To make this book even easier to grasp and read, we compiled three important lists to help you in your study of inorganic chemistry. In Chapter 20, we introduce and explain ten common household products. Then, in Chapter 21, you meet ten of the most important Nobel Prizes that were awarded to chemists. Chapter 22 introduces ten instruments and techniques that are commonly found and used in laboratories across the globe. And finally we give you ten experiments that you can try out at home in Chapter 23. Remember, one of the most fun parts of chemistry is doing chemistry, and this chapter gives you some fun experiments to try.
Icons Used in This Book
Throughout this book icons are used to draw your attention to certain information.
This is not often used here, but the Tip icon indicates that some information may be especially useful to you.
When you see the Remember icon you should understand that this information is quite important to understanding the concepts being explained. If you are studying inorganic chemistry, this is one of the most important icons to look for. It can indicate a definition, or be a concise explanation of a concept; at other times it indicates information to help you grasp how various concepts overlap.
The Warning icon tells you to pay close attention to what’s being said because it indicates where a potentially dangerous situation may arise.
The Technical Stuff icon is used to indicate detailed information; for some people, it might not be necessary to read or understand.
Where to Go from Here
You might be taking an inorganic chemistry course, or maybe you’re just curious about the world around you. Regardless, if you’re looking for something specific, you can find it by checking the index or maybe even the glossary. When you know where to find what you are looking for, go right ahead and jump in. And enjoy.
Part I
Reviewing Some General Chemistry
In this part . . .
You navigate through some of the basic rules of the road that help guide you as you travel through the science of inorganic chemistry. This starts with a definition of inorganic chemistry and continues with a description of the foundation upon which this subject stands. Inorganic chemistry is the study of all the materials known to humankind, and it includes the study of how all the materials interact with one another.
Chapter 1
Introducing Inorganic Chemistry
In This Chapter
Getting familiar with basic concepts in chemistry
Building your knowledge of chemical bonding
Traveling across the periodic table
Delving into details with special topics
Counting by tens: products, prizes, instruments, and experiments
Inorganic chemistry is a practical science. By studying it, you become familiar with the intricate working of processes and materials — from how silicon works in a semiconductor to the reason why steel is stronger than iron. Inorganic chemistry is important for civilization and technological development.
The science of inorganic chemistry covers a great deal of material; in short, it’s the chemistry of everything you see around you. Inorganic chemistry explores and defines laws that atoms follow when they interact, including trends in how they react, characteristics they possess, and the materials they make. It may seem daunting at first to think about how many possibilities there are in the science of inorganic chemistry. Fortunately, each new concept builds on another concept in a very logical way.
This chapter explains what to expect when reading this book and should help you find the right section to guide you through your study of inorganic chemistry.
Building the Foundation
Before diving into the particular details of inorganic chemistry, it’s helpful to understand some of the prominent ideas in general chemistry that are useful to further appreciate inorganic chemistry.
What difference does it make?
It’s important to be able to distinguish between inorganic and organic chemistry. Organic chemistry deals primarily with the reaction of carbon, and its many interactions. But inorganic chemistry deals with all of the other elements (including carbon, too), and it details the various reactions that are possible with each of them. There are a huge number of examples in everyday life that can be described by inorganic chemistry — for example, why metals have so many different colors, or why metal compounds of the same metal can have such varying colors too, like the ones that are used and pigments in paints. It can help to explain how alloys form and what alloys are stronger than others. Or why a dentist uses an acid to open the pores in your teeth before applying an adhesive to make a filling hold fast.
Chemistry is a science of change. It looks at how individual atoms interact with each other and how they are influenced by their environment. We start by explaining what atoms look like, and we describe details of their structure. This is important because the way that the atom is made up determines how reactive that atom is, and as a result of the activity, it can be used by a chemist to make materials. After you have these basics down, you are able to understand the physical properties of many materials based on what atoms they are made from, and why they are made using those specific atoms.
Stemming from this basis of general chemistry we then deal with the specifics of inorganic chemistry. This includes an understanding of approximately 100 atoms that are of practical interest to chemists. To simplify this, inorganic chemistry is understood according to some general trends based on atomic structure that affect the reactivity and bonding of those atoms. This is quite different from the study of organic chemistry that deals with the reactions of just a few atoms, such as carbon, oxygen, nitrogen, and hydrogen. But there is an overlap between inorganic chemistry and organic chemistry in the study of organometallic compounds.
Losing your electrons
In chemical reactions, follow the electrons because electrons hold the key to understanding why reactions take place. Electrons are negatively charged, mobile, and can move from atom to atom; they can be stripped from atoms, too. Atoms are always trying to have just the right amount of electrons to keep stable. If a stable atom has cause to lose or gain an electron, it becomes reactive and starts a chemical process.
The nucleus of an atom has a positive force that attracts electrons. This comes from protons within the nucleus that influence electrons to orbit around the nucleus. As you progress in atomic size, one proton at a time, there is room for one more electron to orbit around the atom.
There are periodic trends that can be seen in the periodic table, the first of which deals with the stability of atoms according to the number of outer electrons in the atom. This is known as valency, and it can be used to show why some atoms are more reactive than others. There are many more periodic trends that are associated with the electrons around the atoms, and you can find more examples in Chapter 2.
Take a stable atom, such as iron, for example. Imagine that you remove an electron from iron; it now has a different reactivity. This is known as oxidation chemistry, and it’s the focus of Chapter 3. The chemistry of oxidation tracks how electrons are gained or lost from molecules, atoms, or ions. When an electron is lost, the molecule, atom, or ion is said to have an increased oxidation state, or is considered oxidized. When the opposite occurs and a molecule, atom, or ion gains an electron, its oxidation state is reduced.
Originally named from the common involvement of oxygen molecules in these types of reactions, chemists now realize that oxidation and reduction reactions (sometimes referred to as redox chemistry) can occur among molecules, atoms, and ions without oxygen.
Splitting atoms: Nuclear chemistry
Another area of general chemistry with which you should be familiar is the study of radioactivity, or nuclear chemistry. Specifically, nuclear chemistry deals with the properties of the nucleus of the atoms; that’s why it is called nuclear chemistry.
As you progress through the periodic table each successive atom has one more proton and neutroncompared with the previous atom. The protons are useful for attracting electrons, and the neutrons are useful for stabilizing the nucleus. When there is an imbalance between the two nuclear particles (proton and neutron), the nucleus becomes unstable, and these types of atoms are called isotopes. If they are radioactive, they are called radioisotopes, and they can be useful, for example, in medical applications.
Although you may immediately think about nuclear reactors for energy, or nuclear bombs and their incredible devastation, concepts in nuclear chemistry are applied for many other, less dramatic purposes, one such example is carbon dating of ancient materials (see Chapter 4).
The nuclear processes can affect the properties of the atoms, and this can have an effect on the properties of materials that are made with those atoms. For example, there is often a lot of heat generated by radioactive atoms, and this heat can affect material properties. Did you know that much of the potassium in our body is in the form of a radioactive isotope? This accounts for some of the heating within our own bodies (see Chapter 11).
Changing pH
In Chapter 5, we explain the basics of acids and bases, including how the pH scale was developed to quantify the strength of different acids and bases. It’s a simple system that ranges in value from pH 1 to pH 14.
Acids have low pH values in the range of pH 1 to pH 7. Bases have high pH values that range from pH 7 to pH 14. In the middle there is pH 7, and this is considered neutral pH, which is also the pH of water. And subsequently is nearly the same pH as blood, demonstrating how important water is to us.
The pH of blood is highly sensitive; if it changes too much, we can get very sick. The preferred range for maintaining stable health is from pH 7.35 to pH 7.45, making blood slightly basic. This simple fact alone highlights the importance of green foods in your diet; they’re alkalizing in your body and help maintain a healthy you.
Chemists have been working for many years to sort out what specifically makes something an acid or a base. Through this work, multiple definitions of acids and bases have been proposed. As we explain in Chapter 5, there are two important models for examining acid-base chemistry:
Brønsted-Lowry model: In this model, an acid is a proton (H) donor, whereas a base accepts hydroxyl groups (OH molecule).
Lewis model: In this model, acids are electron pair acceptors and bases are electron pair donors.
Earlier we said you needed to track the electrons to understand what is happening in various chemical reactions. By using the Lewis model that deals with electron pairs, you can get a good understanding of how reactions occur, by tracking the electron pairs and seeing where they come and go.
It’s important to understand the distinction between these two models. The Brønsted-Lowry model was developed when acids and bases were thought to work in aqueous solvents. As a result, it deals only with hydrogen and hydroxyl groups. On the other hand, the Lewis model was developed to show what happens when water isn’t the solvent, so it deals with electrons instead.
Getting a Grip on Chemical Bonding
Part II delves into how bonding occurs between atoms, and how to distinguish between the types of bonds that are created. Bonding between atoms is important for all scientists to understand because it affects the properties and applications of materials in profound ways. In practice, there are about 100 atoms that are stable enough to form bonds, but there are only three types of bonding known:
Covalent: Covalent bonding stems from the sharing of electrons and the overlap and sharing of electrons orbitals between atoms. Covalent bonds are very strong as a result of this. Covalent bonds have directionality, or a preference for a specific orientation relative to one another, this results in molecules of interesting and specific shapes. As a result, elaborate molecules can be made that have specific structures and symmetry, which we describe in Chapter 7.
Ionic:Ionic bonding occurs when atoms donate or receive electrons rather than share them. One ion is positively charged, and it’s balanced by an ion that is negatively charged; they’re known as the cation and the anion, respectively. Each ion is treated as if it’s a spherical entity with no distortion of the electron orbital. See more information in Chapter 8.
Metallic:Metallicbonds are similar to ionic bonds, so we describe them both in Chapter 8. The main difference is that in metallic bonds the electrons are shared among all the other atoms in the metal materials. This is known as the delocalization of electrons because they are not found locally around one particular atom. This gives rise to many of the properties of metals.
There aren’t strict lines between each type of bond, and sometimes the way atoms bond together is a combination or mixture of more than one bond type. Throughout Part II we explain each of the bond types individually; then in Chapter 9 we will look at how they each influence the formation of molecules known as coordination complexes, which include metallic compounds and connecting molecules called ligands.
Traveling Across the Periodic Table
There are over 100 known atoms, and it can be overwhelming to try to remember each and every one of them. This is what chemists tried to do before the periodictable was created. In Part III, you learn about this important chart that organizes the elements according to their similarities in structure and reactivity. The simplicity and beauty of the periodic table makes it easier to find and compare elements against each other. If the familiar expression “a picture is worth a thousand words” was used to describe inorganic chemistry, then the picture that best describes it is the periodic table. We’ve devoted the chapters in Part III to exploring the periodic table from one end to the other and describing the key characteristics of each group.
Here you can see what the periodic table looks like. Notice how there are 18 groups from left to right as seen at the top. And there are seven periods going from top to bottom as shown on the left side of the table.
Figure 1-1: The periodic table of the elements.
Hyping up hydrogen
Hydrogen is one of the most abundant elements in the universe, and Chapter 10 explains the unique and important properties. This element sits at the upper-left corner of the periodic table and serves as the first step in a long line of stepping stones for you to travel across the periodic table. Some points to know about hydrogen include:
Hydrogen is highly reactive. It lacks one electron in the outer orbital to make it stable, so it has a very reactive valency. This makes it explosive, and for this reason it’s usually found as H2 — two hydrogen atoms bonded together. Because each hydrogen shares the electron, it pacifies the atom.
Hydrogen is used in a technique called nuclear magnetic resonance. This is important because it can be used to elaborate exactly where hydrogen atoms are within a molecule so it can show the structure of the molecule.
Hydrogen can bond with nearly every single atom on the periodic table, making it a versatile atom.
Moving through the main groups
The most common elements are found in the main groups of the periodic table. The main group elements comprise many of the materials we know from everyday experience.
The main group elements include the Group 1 and Group 2 elements on the left side of the table along with Groups 13, 14, 15, 16, 17, and 18 on the right side of the table. The most reactive is on the left side; The most inert and calm reside on the far right. As you might expect, the middle atoms have mixed qualities between these two extremes.
A few of the main group elements have specific qualities recognized by chemists. For example:
Alkali and alkaline earth metals: The elements in the first two columns of the periodic table (excluding hydrogen) are formally known as the alkali and alkaline earth metals, or s-block elements. They are highly reactive and often explosive elements, but also extremely important in biology. Compounds made with Group 1 and 2 elements are often referred to as salts; skip ahead to Chapter 11 to find out why.
Noble gases: The elements in the far right column of the periodic table are the noble gases and are mirror opposites of the alkali and alkaline earth metals. Instead of being reactive, for the most part they are inert, or nonreactive. The noble gases have no need for more electrons, so they generally don’t react with other atoms to gain, give, or share electrons. There are some exceptions, however, because the gases of argon, krypton, and xenon can form compounds with fluorine. More of this can be found in Chapter 12.
The rest of the main group elements, called p-block elements, contain the atoms that are associated with life and living matter, including carbon, oxygen, and nitrogen. More information can be found in Chapters 12 and 17.
Transitioning from one side of the table to another
In the center of the periodic table are the elements that transition from the s-block main group elements to the p-block main group elements. These elements are called the transition metals or d-block elements. The transition metals act as cushion between the highly reactive elements on the far left and the less reactive elements on the right.
These elements are important for industry and help in the synthesis of organic molecules and medicinal compounds. You can find a number of them in the catalytic converter of your car, for example.
Transitional metals are important because they’re used as catalysts in the chemical industry. They’re often reactive atoms, and under the appropriate conditions can complete reactions and make large amounts of molecules with a very specific size and shape. Much of the plastic materials that are in use today are made possible on such a grand and industrial scale thanks to the development of catalysis using transition metals. More information about catalysis can be found in Chapter 16. Catalysts make short work of specificchemical reactions; they have the ability to create a product faster, and with less energy.
Some of the reactions that take place in the body do so because of transition metals. For example, the oxygen that we breathe is carried around the body using a compound that has iron at the center. This is called hemoglobin. But the other transition metals can play important roles in the body also, for more information see Chapter 17.
Many transition metals are used in everyday materials that we use regularly. These metals often have interesting electronic and magnetic properties, and because of this they’re commonly used in electronic devices. But at the nanoscale (that being the very small scale), they have some other very interesting properties that can be harnessed. For more information about nanotechnology, check out Chapter 19.
Uncovering lanthanides and actinides
Buried deep inside the transition metals are two more groups with important, unique characteristics — lanthanides and actinides. They are unique because they use orbital shells that aren’t important to the rest of the periodic table. The chemistry of these materials are not fully understood yet, because some are rare and hard to find, whereas others are radioactive and dangerous to work with. For more information about these elements, see Chapter 14.
Diving Deeper: Special Topics
In Part IV, you get the opportunity to explore some of the more specialized subfields of inorganic chemistry. Each chapter introduces you to how inorganic chemistry is used in a specific way, such as increasing reaction speed (catalysis), or capturing energy from the sunlight (in a chemical reaction called photosynthesis), and building smaller and smaller computer devices. In each chapter, we only brush the surface of these fascinating special topics. But you have enough of the working tools to further your own detailed study of these topics when you want.
Bonding with carbon: Organometallics
In Chapter 14, we introduce the field of organometallic chemistry. As the name suggests, it deals with the chemistry of carbon-containing (or organic) molecules called ligands that bond with metals to form organometallic compounds. Organometallic chemistry combines some aspects of organic chemistry with some aspects of metallic chemistry, and the results are compounds with some unique traits, such as:
The effect of the ligands can be so significant that the colors can be bright blue, red, or green, depending on what ligands are used and where they are placed around the metal center. Atoms with the same metal center can have very bright and brilliant color changes with the addition of different ligands. Many of these compounds are used as pigments in paints.
Most of the organometallic compounds are made with transition metals as the metal center. These metals can have differing magnetic properties depending on the oxidation states, which can be controlled by the placement and type of ligands that are used around the metal.
Organometallic compounds are often used as catalysts. Because they can have very specific geometries, they can make very specific chemical reactions occur.
Speeding things up: Catalysts
Imagine how much more work you could get done if you found a short cut that’s faster and has greater precision in producing results. In chemistry this is possible thanks to catalysis. Catalysis is the chemistry of making things happen faster, or making them happen with less required energy, or both. Catalysis is carried out by chemicals that are called catalysts. A catalyst makes light work out of heavy-duty chemistry. Catalysts are important because they allow for the quick and cheap production of strong and durable materials, such as plastics.
Inside and out: Bio-inorganic and environmental chemistry
You don’t just find examples of inorganic chemistry in the laboratory or in industry; you can also find them inside yourself or around your environment. For instance, the oxygen you’re inhaling right now is being transported around your body by an iron compound inside a large organometallic molecule called hemoglobin. In Chapter 17, we explain how and why this works. Other examples of bio-inorganic chemistry that are described in Chapter 18 include:
Photosynthesis: The chemical reactions involved in photosynthesis transform sunlight energy and carbon dioxide molecules into sugar, water, and oxygen molecules.
Nitrogen fixation: Some bacteria perform chemical reactions that capture atmospheric nitrogen and fix it so that it can be absorbed by organisms (usually plants) through a series of inorganic chemical reactions. The importance of this chemistry can’t be over emphasized. Nitrogen is extremely important to living matter, and nature has developed efficient methods using enzymes in bacteria to work with nitrogen. Science has only recently created similar tools to do so, albeit much more crude than the way that nature does.
Enzymes: Enzymes are proteins that act as catalysts for important functions within your body. Take for example, lactase — the enzyme that’s used to help with the digestion of milk. Some people are lactose intolerant because they lack this enzyme, but they can overcome this by consuming a pill that contains lactase.
Solid-state chemistry
Solid-state chemistry is based on the study of atoms that combine to build solid structures, or crystals. In Chapter 18, you learn how solid-state chemists describe the shape of crystal structures and how this determines the size and shape of the unit cell, which is then used to characterize the many different forms that solid structures take. For example:
Simple crystal structures: Simple crystal structures are composed of atoms that are positioned on the edges of the unit cell.
Binary crystal structures: Binary crystal structures are made of two type of atoms in the crystal, such as NaCl (table salt), for example.
Complex crystal structures: These are more involved than the other examples because they can have more than two different types of atoms present.
One of the most important advances in solid state chemistry is the development of silicon-based materials. The Silicon Valley is where the semiconductor industry was born; scientists worked very hard to learn how to purify silicon and arrange the silicon atoms in such a way that they can be used to make a computer chip. At the heart of every single computer, and most electronic devices, is silicon. Just look around you and imagine a world without silicon, it would be a very different place.
Nanotechnology
In the final chapter of Part IV, we tackle a very new and exciting field called nanotechnology. In this area of study, the size and shape of materials is often of paramount importance. At the size scale of living matter (bacteria are 20 nanometers (nm) in size, DNA is 1-2 nm wide) inorganic chemists can make exquisite materials with near-atomic precision. The advantage of nanotechnology is realized in many different applications; for example, it can be used to enhance catalytic processes, in biomedical applications, and to enhance the mechanical properties of bulk materials.
Nanotechnology is one of the most recent developments to arise from the sciences. It was developed only a couple of decades ago, but already the number of scientific publications and discoveries has been staggering. One of the unique features of this area is that it is important not just for the development of chemistry research, but also physics and biology, too. For this reason many new developments are occurring due to collaborations among researchers of physics, chemistry, and biology. Some of these include foldable electronics, anti-cancer treatments, ever smaller computers, and new methods of water filtration, to name just a few.
Nature has been working at the nanoscale for eons, only now can humankind begin to work at this scale, too. This final chapter gives a brief introduction to the major findings and applications of nanotechnology, but in no way gives full justice to the vast amount of work being carried out in this field. We hope that upon reading Chapter 19 you agree that the future is nano, and that inorganic chemistry plays a vital role in the continued development of this technology.
Listing 40 More
The last part of this book (Part V and the Part of Tens) gives you some nontechnical information about inorganic chemistry. We start right at home by listing some of the common household products that involve inorganic chemical reactions, or inorganic compounds in Chapter 20. These household items may come in handy if you want to try out any of the ten experiments listed in Chapter 23.
Chapter 21 describes ten chemists (or teams of chemists) who have played an important role for inorganic chemistry and who were recognized for their achievements by receiving the Nobel Prize. Finally, Chapter 22 describes ten of the more useful and interesting techniques used in inorganic chemistry research.
Chapter 2
Following the Leader: Atomic Structure and Periodic Trends
In This Chapter
Arranging periodic arrays
Understanding the groupings of elements
Energizing the nucleus
One night, an electron, proton, and neutron came together to form an atom. They were so excited, they decided to go out for a fancy dinner to celebrate. When it came time to pay the tab, the three particles decided to have the bill split evenly amongst themselves. When the waiter returned, he handed the electron and the proton separate receipt books. Confused, the neutron asks: “Where’s mine?” The waiter smirked and said, “For you, sir, there’s no charge.”
This chapter explores how these three critical particles (neutrons, protons, and electrons) render the structures of the numerous atoms we interact with, as well as how repeating trends can be used to predict properties of unknown elements. Atoms are critically important to chemistry. Just as a tower made of Legos can be taken apart to its individual bricks, all of the molecules that make up everything around you can as well. These bricks can then be sorted by all the bricks that have the same properties: You could make a pile of small green pieces, small yellow pieces, and large red pieces, for example. Although a giant tub of Legos can make thousands of different designs, each of those designs stems from the same, limited number of unique pieces. Similarly, the Chemical Abstract Service (CAS) Registry, a list of known organic and inorganic substances, has over 64 million molecules and grows at a rate of 15,000 molecules a day all constructed from nature’s limited number of building blocks — currently scientists have only discovered 118!
Up an’ Atom: Reviewing Atomic Terminology
There are three subatomic particles, or particles smaller than an atom, that comprise the matter in the world around us. Everything we see, touch, smell, taste, and so on, is made of atoms, the basic building blocks of all matter. In turn, each of these atoms contain a combination of:
Neutrons: Neutrally charged particles found in the nucleus of an atom.
Protons: Positively charged particles, also found in the nucleus; it’s important to note that the number of protons an atom possesses is the sole factor that distinguishes one element from another.
Electrons: Negatively charged particles not found in the nucleus, but at the core of most all chemical reactions.
It’s important to remember that an element is defined by the number of protons it has. For example, all carbon atoms, by definition, have six protons; however, many isotopes, or atoms with the exact same number of protons but different numbers of neutrons, of carbon exist. The most common three isotopes of carbon are carbon-12, carbon-13, and carbon-14. Each of these isotopes has six protons, yet a varying number of neutrons (six, seven, and eight, respectively). The sum of the protons and neutrons make up the mass number.
Figure 2-1 shows the nuclear notation (a way of writing elements that gives information about the nucleus of element) for the polyatomic ion with 16 neutrons on each sulfur atom, or nuclide.
Figure 2-1: Nuclear notation for polyatomic ion with 16 neutrons on each sulfur atom.
Finding atoms
The story of the atom is a wonderful tale that has existed for centuries, before scientists could even prove that an atom, in fact, exists.
Around 450 B.C.: Leucippus and his pupil Democritus develop the idea that all matter is made up of atoms, an ancient Greek word for indivisible, using the logic argument of everything can only be divided a finite number of times until it is too small to be further divided. At that point, everything is atoms (indivisible particles) and empty space. Moreover, the atoms must take on properties of their bulk materials: strong iron must have hooks that hold the iron atoms together, while water must be smooth to allow it to flow, for example. However, this theory took a backseatto Aristotle's classical five elements: earth, wind, water, air, and aether for many centuries.
1803 A.D.: John Dalton notices that reactions occur in specific proportions based on the respective weights and determines the relative weights of six elements (H, O, N, C, S, and P). From this, Dalton created the first points of modern atomic theory:
Elements are made of atoms.
Atoms of any element are identical.
Atoms cannot be subdivided, created, or destroyed.
Atoms form compounds in whole-number ratios.
Chemical reactions are the rearrangement of atoms.
Although many of these postulates were later disproven, these underlie a major turning point in the chemist’s view of the world, earning Dalton the title as one of the fathers of modern chemistry.
1897 – 1904 A.D.: Following the recently discovered negatively-charged electron (called “corpuscles by their discoverer J. J. Thompson), the plum pudding model of the atom was born. In this model, the atom remained a definite shape, like Democritus and Leucippus proposed, but with the majority of the space being positively charged with tiny specks of negatively-charged particles distributed throughout like plums throughout plum pudding (or like chocolate chips throughout a chocolate chip cookie).
1909 A.D.: Rutherford and his team of scientists developed the idea of a positively-charged nucleus surrounded by mostly empty space after bombarding many metal foils, most famously his gold foil with alpha particles, making the curious observation that a tiny amount of the alpha particles bounced back. Rutherford famously described this to be as amazing as shooting a cannonball at a piece of tissue paper and having the cannonball bounce back. The only explanation was a highly concentrated nucleus of positive charge at the center of an atom.
1913 A.D.: A Danish physicist named Niels Bohr envisaged electrons orbiting the nucleus in discrete orbits, where each orbit has a very specific energy level that could explain the different (but very specific) energies of light that were emitted from different elements. Essentially, this model mimics the planets of our solar system orbiting our sun, but relies on the force of attraction between the negative electrons and the positive nucleus instead of gravity like our solar system uses.
1950s A.D.: Quantum mechanics jazzes up all of these theories of the atom by showing that just as traditional wave-like light could be considered a particle, traditional particle-like electrons could be considered like a wave. This wave-particle duality