Chemistry Textbook PDF

Chemistry OpenStax College Rice University 6100 Main Street MS-375 Houston, Texas 77005 To learn more about OpenStax College, visit http://openstaxcollege.org. Individual print copies and bulk orders can be purchased through our website. © 2015 Rice University. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution 4.0 International License. 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Table of Contents Preface................................................ 1 Chapter 1: Essential Ideas...................................... 9 1.1 Chemistry in Context................................... 10 1.2 Phases and Classification of Matter........................... 15 1.3 Physical and Chemical Properties............................ 25 1.4 Measurements...................................... 29 1.5 Measurement Uncertainty, Accuracy, and Precision................... 36 1.6 Mathematical Treatment of Measurement Results.................... 44 Chapter 2: Atoms, Molecules, and Ions.............................. 67 2.1 Early Ideas in Atomic Theory............................... 68 2.2 Evolution of Atomic Theory................................ 72 2.3 Atomic Structure and Symbolism............................. 78 2.4 Chemical Formulas.................................... 87 2.5 The Periodic Table.................................... 93 2.6 Molecular and Ionic Compounds............................. 97 2.7 Chemical Nomenclature................................. 105 Chapter 3: Composition of Substances and Solutions....................... 129 3.1 Formula Mass and the Mole Concept.......................... 130 3.2 Determining Empirical and Molecular Formulas..................... 142 3.3 Molarity.......................................... 149 3.4 Other Units for Solution Concentrations......................... 157 Chapter 4: Stoichiometry of Chemical Reactions......................... 175 4.1 Writing and Balancing Chemical Equations....................... 176 4.2 Classifying Chemical Reactions............................. 182 4.3 Reaction Stoichiometry.................................. 196 4.4 Reaction Yields...................................... 201 4.5 Quantitative Chemical Analysis.............................. 206 Chapter 5: Thermochemistry.................................... 231 5.1 Energy Basics...................................... 232 5.2 Calorimetry........................................ 242 5.3 Enthalpy......................................... 255 Chapter 6: Electronic Structure and Periodic Properties of Elements............... 281 6.1 Electromagnetic Energy................................. 282 6.2 The Bohr Model...................................... 296 6.3 Development of Quantum Theory............................ 301 6.4 Electronic Structure of Atoms (Electron Configurations)................. 315 6.5 Periodic Variations in Element Properties........................ 324 Chapter 7: Chemical Bonding and Molecular Geometry...................... 345 7.1 Ionic Bonding....................................... 346 7.2 Covalent Bonding..................................... 349 7.3 Lewis Symbols and Structures.............................. 355 7.4 Formal Charges and Resonance............................. 365 7.5 Strengths of Ionic and Covalent Bonds.......................... 370 7.6 Molecular Structure and Polarity............................. 377 Chapter 8: Advanced Theories of Covalent Bonding........................ 413 8.1 Valence Bond Theory................................... 414 8.2 Hybrid Atomic Orbitals.................................. 418 8.3 Multiple Bonds...................................... 430 8.4 Molecular Orbital Theory................................. 433 Chapter 9: Gases.......................................... 459 9.1 Gas Pressure....................................... 460 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law....... 469 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions............ 483 9.4 Effusion and Diffusion of Gases............................. 495 9.5 The Kinetic-Molecular Theory.............................. 500 9.6 Non-Ideal Gas Behavior................................. 506 Chapter 10: Liquids and Solids................................... 525 10.1 Intermolecular Forces.................................. 526 10.2 Properties of Liquids.................................. 539 10.3 Phase Transitions.................................... 545 10.4 Phase Diagrams..................................... 556 10.5 The Solid State of Matter................................ 564 10.6 Lattice Structures in Crystalline Solids......................... 570 Chapter 11: Solutions and Colloids................................. 603 11.1 The Dissolution Process................................. 604 11.2 Electrolytes....................................... 609 11.3 Solubility......................................... 612 11.4 Colligative Properties.................................. 622 11.5 Colloids......................................... 641 Chapter 12: Kinetics........................................ 661 12.1 Chemical Reaction Rates................................ 662 12.2 Factors Affecting Reaction Rates............................ 667 12.3 Rate Laws........................................ 670 12.4 Integrated Rate Laws.................................. 678 12.5 Collision Theory..................................... 689 12.6 Reaction Mechanisms.................................. 695 12.7 Catalysis......................................... 701 Chapter 13: Fundamental Equilibrium Concepts.......................... 729 13.1 Chemical Equilibria................................... 730 13.2 Equilibrium Constants.................................. 734 13.3 Shifting Equilibria: Le Châtelier’s Principle....................... 743 13.4 Equilibrium Calculations................................. 749 Chapter 14: Acid-Base Equilibria.................................. 777 14.1 Brønsted-Lowry Acids and Bases............................ 778 14.2 pH and pOH....................................... 782 14.3 Relative Strengths of Acids and Bases......................... 789 14.4 Hydrolysis of Salt Solutions............................... 809 14.5 Polyprotic Acids..................................... 817 14.6 Buffers.......................................... 821 14.7 Acid-Base Titrations................................... 829 Chapter 15: Equilibria of Other Reaction Classes......................... 855 15.1 Precipitation and Dissolution.............................. 856 15.2 Lewis Acids and Bases................................. 873 15.3 Multiple Equilibria.................................... 878 Chapter 16: Thermodynamics................................... 901 16.1 Spontaneity....................................... 901 16.2 Entropy......................................... 905 16.3 The Second and Third Laws of Thermodynamics.................... 911 16.4 Free Energy....................................... 916 Chapter 17: Electrochemistry................................... 937 17.1 Balancing Oxidation-Reduction Reactions....................... 938 17.2 Galvanic Cells...................................... 945 This content is available for free at https://cnx.org/content/col11760/1.9 17.3 Standard Reduction Potentials............................. 950 17.4 The Nernst Equation.................................. 956 17.5 Batteries and Fuel Cells................................. 960 17.6 Corrosion........................................ 967 17.7 Electrolysis....................................... 970 Chapter 18: Representative Metals, Metalloids, and Nonmetals.................. 987 18.1 Periodicity........................................ 988 18.2 Occurrence and Preparation of the Representative Metals............... 998 18.3 Structure and General Properties of the Metalloids.................. 1002 18.4 Structure and General Properties of the Nonmetals................. 1011 18.5 Occurrence, Preparation, and Compounds of Hydrogen............... 1019 18.6 Occurrence, Preparation, and Properties of Carbonates............... 1026 18.7 Occurrence, Preparation, and Properties of Nitrogen................. 1029 18.8 Occurrence, Preparation, and Properties of Phosphorus............... 1034 18.9 Occurrence, Preparation, and Compounds of Oxygen................ 1036 18.10 Occurrence, Preparation, and Properties of Sulfur................. 1052 18.11 Occurrence, Preparation, and Properties of Halogens................ 1054 18.12 Occurrence, Preparation, and Properties of the Noble Gases............ 1060 Chapter 19: Transition Metals and Coordination Chemistry................... 1077 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds 1078 19.2 Coordination Chemistry of Transition Metals..................... 1092 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds........ 1107 Chapter 20: Organic Chemistry................................. 1125 20.1 Hydrocarbons..................................... 1126 20.2 Alcohols and Ethers.................................. 1144 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters.................. 1149 20.4 Amines and Amides.................................. 1154 Chapter 21: Nuclear Chemistry................................. 1175 21.1 Nuclear Structure and Stability............................ 1176 21.2 Nuclear Equations................................... 1183 21.3 Radioactive Decay.................................. 1186 21.4 Transmutation and Nuclear Energy.......................... 1198 21.5 Uses of Radioisotopes................................ 1213 21.6 Biological Effects of Radiation............................. 1218 A The Periodic Table....................................... 1239 B Essential Mathematics..................................... 1241 C Units and Conversion Factors................................. 1249 D Fundamental Physical Constants............................... 1251 E Water Properties........................................ 1253 F Composition of Commercial Acids and Bases......................... 1259 G Standard Thermodynamic Properties for Selected Substances................ 1261 H Ionization Constants of Weak Acids.............................. 1277 I Ionization Constants of Weak Bases.............................. 1281 J Solubility Products....................................... 1283 K Formation Constants for Complex Ions............................. 1289 L Standard Electrode (Half-Cell) Potentials............................ 1291 M Half-Lives for Several Radioactive Isotopes.......................... 1299 Index............................................... 1377 This content is available for free at https://cnx.org/content/col11760/1.9 Preface 1 Preface Welcome to Chemistry, an OpenStax College resource. This textbook has been created with several goals in mind: accessibility, customization, and student engagement—all while encouraging students toward high levels of academic scholarship. Instructors and students alike will find that this textbook offers a strong foundation in chemistry in an accessible format. About OpenStax College OpenStax College is a non-profit organization committed to improving student access to quality learning materials. Our free textbooks go through a rigorous editorial publishing process. Our texts are developed and peer-reviewed by educators to ensure they are readable, accurate, and meet the scope and sequence requirements of today’s college courses. Unlike traditional textbooks, OpenStax College resources live online and are owned by the community of educators using them. Through our partnerships with companies and foundations committed to reducing costs for students, OpenStax College is working to improve access to higher education for all. OpenStax College is an initiative of Rice University and is made possible through the generous support of several philanthropic foundations. Since our launch in 2012 our texts have been used by millions of learners online and over 1,091 institutions worldwide. About OpenStax College’s Resources OpenStax College resources provide quality academic instruction. Three key features set our materials apart from others: they can be customized by instructors for each class, they are a "living" resource that grows online through contributions from educators, and they are available free or for minimal cost. Customization OpenStax College learning resources are designed to be customized for each course. Our textbooks provide a solid foundation on which instructors can build, and our resources are conceived and written with flexibility in mind. Instructors can select the sections most relevant to their curricula and create a textbook that speaks directly to the needs of their classes and student body. Teachers are encouraged to expand on existing examples by adding unique context via geographically localized applications and topical connections. Chemistry can be easily customized using our online platform (http://cnx.org/content/col11760/latest). Simply select the content most relevant to your current semester and create a textbook that speaks directly to the needs of your class. Chemistry is organized as a collection of sections that can be rearranged, modified, and enhanced through localized examples or to incorporate a specific theme of your course. This customization feature will ensure that your textbook truly reflects the goals of your course. Curation To broaden access and encourage community curation, Chemistry is “open source” licensed under a Creative Commons Attribution (CC-BY) license. The academic science community is invited to submit examples, emerging research, and other feedback to enhance and strengthen the material and keep it current and relevant for today’s students. Submit your suggestions to [email protected], and check in on edition status, alternate versions, errata, and news on the StaxDash at http://openstaxcollege.org. Cost Our textbooks are available for free online, and in low-cost print and e-book editions. About Chemistry Chemistry is designed for the two-semester general chemistry course. For many students, this course provides the foundation to a career in chemistry, while for others, this may be their only college-level science course. As such, this textbook provides an important opportunity for students to learn the core concepts of chemistry and understand how those concepts apply to their lives and the world around them. The text has been developed to meet the scope and sequence of most general chemistry courses. At the same time, the book includes a number of innovative features designed to enhance student learning. A strength of Chemistry is that instructors can customize the book, adapting it to the approach that works best in their classroom. Coverage and Scope 2 Preface Our Chemistry textbook adheres to the scope and sequence of most general chemistry courses nationwide. We strive to make chemistry, as a discipline, interesting and accessible to students. With this objective in mind, the content of this textbook has been developed and arranged to provide a logical progression from fundamental to more advanced concepts of chemical science. Topics are introduced within the context of familiar experiences whenever possible, treated with an appropriate rigor to satisfy the intellect of the learner, and reinforced in subsequent discussions of related content. The organization and pedagogical features were developed and vetted with feedback from chemistry educators dedicated to the project. Chapter 1: Essential Ideas Chapter 2: Atoms, Molecules, and Ions Chapter 3: Composition of Substances and Solutions Chapter 4: Stoichiometry of Chemical Reactions Chapter 5: Thermochemistry Chapter 6: Electronic Structures and Periodic Properties of Elements Chapter 7: Chemical Bonding and Molecular Geometry Chapter 8: Advanced Theories of Covalent Bonding Chapter 9: Gases Chapter 10: Liquids and Solids Chapter 11: Solutions and Colloids Chapter 12: Kinetics Chapter 13: Fundamental Equilibrium Concepts Chapter 14: Acid-Base Equilibria Chapter 15: Equilibria of Other Reaction Classes Chapter 16: Thermodynamics Chapter 17: Electrochemistry Chapter 18: Representative Metals, Metalloids, and Nonmetals Chapter 19: Transition Metals and Coordination Chemistry Chapter 20: Organic Chemistry Chapter 21: Nuclear Chemistry Pedagogical Foundation Throughout Chemistry, you will find features that draw the students into scientific inquiry by taking selected topics a step further. Students and educators alike will appreciate discussions in these feature boxes. Chemistry in Everyday Life ties chemistry concepts to everyday issues and real-world applications of science that students encounter in their lives. Topics include cell phones, solar thermal energy power plants, plastics recycling, and measuring blood pressure. How Sciences Interconnect feature boxes discuss chemistry in context of its interconnectedness with other scientific disciplines. Topics include neurotransmitters, greenhouse gases and climate change, and proteins and enzymes. Portrait of a Chemist features present a short bio and an introduction to the work of prominent figures from history and present day so that students can see the “face” of contributors in this field as well as science in action. Comprehensive Art Program This content is available for free at https://cnx.org/content/col11760/1.9 Preface 3 Our art program is designed to enhance students’ understanding of concepts through clear, effective illustrations, diagrams, and photographs. 4 Preface Interactives That Engage Chemistry incorporates links to relevant interactive exercises and animations that help bring topics to life through our Link to Learning feature. Examples include: PhET simulations This content is available for free at https://cnx.org/content/col11760/1.9 Preface 5 IUPAC data and interactives TED talks Assessments That Reinforce Key Concepts In-chapter Examples walk students through problems by posing a question, stepping out a solution, and then asking students to practice the skill with a “Check Your Learning” component. The book also includes assessments at the end of each chapter so students can apply what they’ve learned through practice problems. Atom-First Alternate Sequencing Chemistry was conceived and written to fit a particular topical sequence, but it can be used flexibly to accommodate other course structures. Some instructors prefer to organize their course in a molecule-first or atom-first organization. For professors who use this approach, our OpenStax Chemistry textbook can be sequenced to fit this pedagogy. Please consider, however, that the chapters were not written to be completely independent, and that the proposed alternate sequence should be carefully considered for student preparation and textual consistency. We recommend these shifts in the table of contents structure if you plan to create a molecule/atom-first version of this text for your students: Chapter 1: Essential Ideas Chapter 2: Atoms, Molecules, and Ions Chapter 6: Electronic Structure and Periodic Properties of Elements Chapter 7: Chemical Bonding and Molecular Geometry Chapter 8: Advanced Theories of Covalent Bonding Chapter 3: Composition of Substances and Solutions Chapter 4: Stoichiometry of Chemical Reactions Chapter 5: Thermochemistry Chapter 9: Gases Chapter 10: Liquids and Solids Chapter 11: Solutions and Colloids Chapter 12: Kinetics Chapter 13: Fundamental Equilibrium Concepts Chapter 14: Acid-Base Equilibria Chapter 15: Equilibria of Other Reaction Classes Chapter 16: Thermodynamics Chapter 17: Electrochemistry Chapter 18: Representative Metals, Metalloids, and Nonmetals Chapter 19: Transition Metals and Coordination Chemistry Chapter 20: Organic Chemistry Chapter 21: Nuclear Chemistry Ancillaries OpenStax projects offer an array of ancillaries for students and instructors. The following resources are available. PowerPoint Slides Instructor’s Solution Manual Our resources are continually expanding, so please visit http://openstaxcollege.org to view an up-to-date list of the Learning Resources for this title and to find information on accessing these resources. 6 Preface About Our Team Content Leads Paul Flowers, PhD, University of North Carolina - Pembroke Dr. Paul Flowers earned a BS in Chemistry from St. Andrews Presbyterian College in 1983 and a PhD in Analytical Chemistry from the University of Tennessee in 1988. After a one-year postdoctoral appointment at Los Alamos National Laboratory, he joined the University of North Carolina–Pembroke in the fall of 1989. Dr. Flowers teaches courses in general and analytical chemistry, and conducts experimental research involving the development of new devices and methods for microscale chemical analysis. Klaus Theopold, PhD, University of Delaware Dr. Klaus Theopold (born in Berlin, Germany) received his Vordiplom from the Universität Hamburg in 1977. He then decided to pursue his graduate studies in the United States, where he received his PhD in inorganic chemistry from UC Berkeley in 1982. After a year of postdoctoral research at MIT, he joined the faculty at Cornell University. In 1990, he moved to the University of Delaware, where he is a Professor in the Department of Chemistry and Biochemistry and serves as an Associate Director of the University’s Center for Catalytic Science and Technology. Dr. Theopold regularly teaches graduate courses in inorganic and organometallic chemistry as well as General Chemistry. Richard Langley, PhD, Stephen F. Austin State University Dr. Richard Langley earned BS degrees in Chemistry and Mineralogy from Miami University of Ohio in the early 1970s and went on to receive his PhD in Chemistry from the University of Nebraska in 1977. After a postdoctoral fellowship at the Arizona State University Center for Solid State Studies, Dr. Langley taught in the University of Wisconsin system and participated in research at Argonne National Laboratory. Moving to Stephen F. Austin State University in 1982, Dr. Langley today serves as Professor of Chemistry. His areas of specialization are solid state chemistry, synthetic inorganic chemistry, fluorine chemistry, and chemical education. Senior Contributing Author William R. Robinson, PhD Contributors Mark Blaser, Shasta College Simon Bott, University of Houston Donald Carpenetti, Craven Community College Andrew Eklund, Alfred University Emad El-Giar, University of Louisiana at Monroe Don Frantz, Wilfrid Laurier University Paul Hooker, Westminster College Jennifer Look, Mercer University George Kaminski, Worcester Polytechnic Institute Carol Martinez, Central New Mexico Community College Troy Milliken, Jackson State University Vicki Moravec, Trine University Jason Powell, Ferrum College Thomas Sorensen, University of Wisconsin–Milwaukee Allison Soult, University of Kentucky Reviewers Casey Akin, College Station Independent School District Lara AL-Hariri, University of Massachusetts–Amherst Sahar Atwa, University of Louisiana at Monroe Todd Austell, University of North Carolina–Chapel Hill Bobby Bailey, University of Maryland–University College Robert Baker, Trinity College This content is available for free at https://cnx.org/content/col11760/1.9 Preface 7 Jeffrey Bartz, Kalamazoo College Greg Baxley, Cuesta College Ashley Beasley Green, National Institute of Standards and Technology Patricia Bianconi, University of Massachusetts Lisa Blank, Lyme Central School District Daniel Branan, Colorado Community College System Dorian Canelas, Duke University Emmanuel Chang, York College Carolyn Collins, College of Southern Nevada Colleen Craig, University of Washington Yasmine Daniels, Montgomery College–Germantown Patricia Dockham, Grand Rapids Community College Erick Fuoco, Richard J. Daley College Andrea Geyer, University of Saint Francis Daniel Goebbert, University of Alabama John Goodwin, Coastal Carolina University Stephanie Gould, Austin College Patrick Holt, Bellarmine University Kevin Kolack, Queensborough Community College Amy Kovach, Roberts Wesleyan College Judit Kovacs Beagle, University of Dayton Krzysztof Kuczera, University of Kansas Marcus Lay, University of Georgia Pamela Lord, University of Saint Francis Oleg Maksimov, Excelsior College John Matson, Virginia Tech Katrina Miranda, University of Arizona Douglas Mulford, Emory University Mark Ott, Jackson College Adrienne Oxley, Columbia College Richard Pennington, Georgia Gwinnett College Rodney Powell, Coastal Carolina Community College Jeanita Pritchett, Montgomery College–Rockville Aheda Saber, University of Illinois at Chicago Raymond Sadeghi, University of Texas at San Antonio Nirmala Shankar, Rutgers University Jonathan Smith, Temple University Bryan Spiegelberg, Rider University Ron Sternfels, Roane State Community College Cynthia Strong, Cornell College Kris Varazo, Francis Marion University Victor Vilchiz, Virginia State University Alex Waterson, Vanderbilt University JuchaoYan, Eastern New Mexico University Mustafa Yatin, Salem State University Kazushige Yokoyama, State University of New York at Geneseo Curtis Zaleski, Shippensburg University Wei Zhang, University of Colorado–Boulder 8 Preface This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 9 Chapter 1 Essential Ideas Figure 1.1 Chemical substances and processes are essential for our existence, providing sustenance, keeping us clean and healthy, fabricating electronic devices, enabling transportation, and much more. (credit “left”: modification of work by “vxla”/Flickr; credit “left middle”: modification of work by “the Italian voice”/Flickr; credit “right middle”: modification of work by Jason Trim; credit “right”: modification of work by “gosheshe”/Flickr) Chapter Outline 1.1 Chemistry in Context 1.2 Phases and Classification of Matter 1.3 Physical and Chemical Properties 1.4 Measurements 1.5 Measurement Uncertainty, Accuracy, and Precision 1.6 Mathematical Treatment of Measurement Results Introduction Your alarm goes off and, after hitting “snooze” once or twice, you pry yourself out of bed. You make a cup of coffee to help you get going, and then you shower, get dressed, eat breakfast, and check your phone for messages. On your way to school, you stop to fill your car’s gas tank, almost making you late for the first day of chemistry class. As you find a seat in the classroom, you read the question projected on the screen: “Welcome to class! Why should we study chemistry?” Do you have an answer? You may be studying chemistry because it fulfills an academic requirement, but if you consider your daily activities, you might find chemistry interesting for other reasons. Most everything you do and encounter during your day involves chemistry. Making coffee, cooking eggs, and toasting bread involve chemistry. The products you use—like soap and shampoo, the fabrics you wear, the electronics that keep you connected to your world, the gasoline that propels your car—all of these and more involve chemical substances and processes. Whether you are aware or not, chemistry is part of your everyday world. In this course, you will learn many of the essential principles underlying the chemistry of modern-day life. 10 Chapter 1 | Essential Ideas 1.1 Chemistry in Context By the end of this module, you will be able to: Outline the historical development of chemistry Provide examples of the importance of chemistry in everyday life Describe the scientific method Differentiate among hypotheses, theories, and laws Provide examples illustrating macroscopic, microscopic, and symbolic domains Throughout human history, people have tried to convert matter into more useful forms. Our Stone Age ancestors chipped pieces of flint into useful tools and carved wood into statues and toys. These endeavors involved changing the shape of a substance without changing the substance itself. But as our knowledge increased, humans began to change the composition of the substances as well—clay was converted into pottery, hides were cured to make garments, copper ores were transformed into copper tools and weapons, and grain was made into bread. Humans began to practice chemistry when they learned to control fire and use it to cook, make pottery, and smelt metals. Subsequently, they began to separate and use specific components of matter. A variety of drugs such as aloe, myrrh, and opium were isolated from plants. Dyes, such as indigo and Tyrian purple, were extracted from plant and animal matter. Metals were combined to form alloys—for example, copper and tin were mixed together to make brass—and more elaborate smelting techniques produced iron. Alkalis were extracted from ashes, and soaps were prepared by combining these alkalis with fats. Alcohol was produced by fermentation and purified by distillation. Attempts to understand the behavior of matter extend back for more than 2500 years. As early as the sixth century BC, Greek philosophers discussed a system in which water was the basis of all things. You may have heard of the Greek postulate that matter consists of four elements: earth, air, fire, and water. Subsequently, an amalgamation of chemical technologies and philosophical speculations were spread from Egypt, China, and the eastern Mediterranean by alchemists, who endeavored to transform “base metals” such as lead into “noble metals” like gold, and to create elixirs to cure disease and extend life (Figure 1.2). This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 11 Figure 1.2 This portrayal shows an alchemist’s workshop circa 1580. Although alchemy made some useful contributions to how to manipulate matter, it was not scientific by modern standards. (credit: Chemical Heritage Foundation) From alchemy came the historical progressions that led to modern chemistry: the isolation of drugs from natural sources, metallurgy, and the dye industry. Today, chemistry continues to deepen our understanding and improve our ability to harness and control the behavior of matter. This effort has been so successful that many people do not realize either the central position of chemistry among the sciences or the importance and universality of chemistry in daily life. Chemistry: The Central Science Chemistry is sometimes referred to as “the central science” due to its interconnectedness with a vast array of other STEM disciplines (STEM stands for areas of study in the science, technology, engineering, and math fields). Chemistry and the language of chemists play vital roles in biology, medicine, materials science, forensics, environmental science, and many other fields (Figure 1.3). The basic principles of physics are essential for understanding many aspects of chemistry, and there is extensive overlap between many subdisciplines within the two fields, such as chemical physics and nuclear chemistry. Mathematics, computer science, and information theory provide important tools that help us calculate, interpret, describe, and generally make sense of the chemical world. Biology and chemistry converge in biochemistry, which is crucial to understanding the many complex factors and processes that keep living organisms (such as us) alive. Chemical engineering, materials science, and nanotechnology combine chemical principles and empirical findings to produce useful substances, ranging from gasoline to fabrics to electronics. Agriculture, food science, veterinary science, and brewing and wine making help provide sustenance in the form of food and drink to the world’s population. Medicine, pharmacology, biotechnology, and botany identify and produce substances that help keep us healthy. Environmental science, geology, oceanography, and atmospheric science incorporate many chemical ideas to help us better understand and protect our physical world. Chemical ideas are used to help understand the universe in astronomy and cosmology. 12 Chapter 1 | Essential Ideas Figure 1.3 Knowledge of chemistry is central to understanding a wide range of scientific disciplines. This diagram shows just some of the interrelationships between chemistry and other fields. What are some changes in matter that are essential to daily life? Digesting and assimilating food, synthesizing polymers that are used to make clothing, containers, cookware, and credit cards, and refining crude oil into gasoline and other products are just a few examples. As you proceed through this course, you will discover many different examples of changes in the composition and structure of matter, how to classify these changes and how they occurred, their causes, the changes in energy that accompany them, and the principles and laws involved. As you learn about these things, you will be learning chemistry, the study of the composition, properties, and interactions of matter. The practice of chemistry is not limited to chemistry books or laboratories: It happens whenever someone is involved in changes in matter or in conditions that may lead to such changes. The Scientific Method Chemistry is a science based on observation and experimentation. Doing chemistry involves attempting to answer questions and explain observations in terms of the laws and theories of chemistry, using procedures that are accepted by the scientific community. There is no single route to answering a question or explaining an observation, but there is an aspect common to every approach: Each uses knowledge based on experiments that can be reproduced to verify the results. Some routes involve a hypothesis, a tentative explanation of observations that acts as a guide for gathering and checking information. We test a hypothesis by experimentation, calculation, and/or comparison with the experiments of others and then refine it as needed. Some hypotheses are attempts to explain the behavior that is summarized in laws. The laws of science summarize a vast number of experimental observations, and describe or predict some facet of the natural world. If such a hypothesis turns out to be capable of explaining a large body of experimental data, it can reach the status of a theory. Scientific theories are well-substantiated, comprehensive, testable explanations of particular aspects of nature. Theories are accepted because they provide satisfactory explanations, but they can be modified if new data become available. The path of discovery that leads from question and observation to law or hypothesis to theory, combined with experimental verification of the hypothesis and any necessary modification of the theory, is called the scientific method (Figure 1.4). This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 13 Figure 1.4 The scientific method follows a process similar to the one shown in this diagram. All the key components are shown, in roughly the right order. Scientific progress is seldom neat and clean: It requires open inquiry and the reworking of questions and ideas in response to findings. The Domains of Chemistry Chemists study and describe the behavior of matter and energy in three different domains: macroscopic, microscopic, and symbolic. These domains provide different ways of considering and describing chemical behavior. Macro is a Greek word that means “large.” The macroscopic domain is familiar to us: It is the realm of everyday things that are large enough to be sensed directly by human sight or touch. In daily life, this includes the food you eat and the breeze you feel on your face. The macroscopic domain includes everyday and laboratory chemistry, where we observe and measure physical and chemical properties, or changes such as density, solubility, and flammability. The microscopic domain of chemistry is almost always visited in the imagination. Micro also comes from Greek and means “small.” Some aspects of the microscopic domains are visible through a microscope, such as a magnified image of graphite or bacteria. Viruses, for instance, are too small to be seen with the naked eye, but when we’re suffering from a cold, we’re reminded of how real they are. However, most of the subjects in the microscopic domain of chemistry—such as atoms and molecules—are too small to be seen even with standard microscopes and often must be pictured in the mind. Other components of the microscopic domain include ions and electrons, protons and neutrons, and chemical bonds, each of which is far too small to see. This domain includes the individual metal atoms in a wire, the ions that compose a salt crystal, the changes in individual molecules that result in a color change, the conversion of nutrient molecules into tissue and energy, and the evolution of heat as bonds that hold atoms together are created. The symbolic domain contains the specialized language used to represent components of the macroscopic and microscopic domains. Chemical symbols (such as those used in the periodic table), chemical formulas, and chemical equations are part of the symbolic domain, as are graphs and drawings. We can also consider calculations as part of the symbolic domain. These symbols play an important role in chemistry because they help interpret the behavior of the macroscopic domain in terms of the components of the microscopic domain. One of the challenges for 14 Chapter 1 | Essential Ideas students learning chemistry is recognizing that the same symbols can represent different things in the macroscopic and microscopic domains, and one of the features that makes chemistry fascinating is the use of a domain that must be imagined to explain behavior in a domain that can be observed. A helpful way to understand the three domains is via the essential and ubiquitous substance of water. That water is a liquid at moderate temperatures, will freeze to form a solid at lower temperatures, and boil to form a gas at higher temperatures (Figure 1.5) are macroscopic observations. But some properties of water fall into the microscopic domain—what we cannot observe with the naked eye. The description of water as comprised of two hydrogen atoms and one oxygen atom, and the explanation of freezing and boiling in terms of attractions between these molecules, is within the microscopic arena. The formula H2O, which can describe water at either the macroscopic or microscopic levels, is an example of the symbolic domain. The abbreviations (g) for gas, (s) for solid, and (l) for liquid are also symbolic. Figure 1.5 (a) Moisture in the air, icebergs, and the ocean represent water in the macroscopic domain. (b) At the molecular level (microscopic domain), gas molecules are far apart and disorganized, solid water molecules are close together and organized, and liquid molecules are close together and disorganized. (c) The formula H2O symbolizes water, and (g), (s), and (l) symbolize its phases. Note that clouds are actually comprised of either very small liquid water droplets or solid water crystals; gaseous water in our atmosphere is not visible to the naked eye, although it may be sensed as humidity. (credit a: modification of work by “Gorkaazk”/Wikimedia Commons) This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 15 1.2 Phases and Classification of Matter By the end of this section, you will be able to: Describe the basic properties of each physical state of matter: solid, liquid, and gas Define and give examples of atoms and molecules Classify matter as an element, compound, homogeneous mixture, or heterogeneous mixture with regard to its physical state and composition Distinguish between mass and weight Apply the law of conservation of matter Matter is defined as anything that occupies space and has mass, and it is all around us. Solids and liquids are more obviously matter: We can see that they take up space, and their weight tells us that they have mass. Gases are also matter; if gases did not take up space, a balloon would stay collapsed rather than inflate when filled with gas. Solids, liquids, and gases are the three states of matter commonly found on earth (Figure 1.6). A solid is rigid and possesses a definite shape. A liquid flows and takes the shape of a container, except that it forms a flat or slightly curved upper surface when acted upon by gravity. (In zero gravity, liquids assume a spherical shape.) Both liquid and solid samples have volumes that are very nearly independent of pressure. A gas takes both the shape and volume of its container. Figure 1.6 The three most common states or phases of matter are solid, liquid, and gas. A fourth state of matter, plasma, occurs naturally in the interiors of stars. A plasma is a gaseous state of matter that contains appreciable numbers of electrically charged particles (Figure 1.7). The presence of these charged particles imparts unique properties to plasmas that justify their classification as a state of matter distinct from gases. In addition to stars, plasmas are found in some other high-temperature environments (both natural and man-made), such as lightning strikes, certain television screens, and specialized analytical instruments used to detect trace amounts of metals. 16 Chapter 1 | Essential Ideas Figure 1.7 A plasma torch can be used to cut metal. (credit: “Hypertherm”/Wikimedia Commons) Link to Learning In a tiny cell in a plasma television, the plasma emits ultraviolet light, which in turn causes the display at that location to appear a specific color. The composite of these tiny dots of color makes up the image that you see. Watch this video (http://openstaxcollege.org/l/16plasma) to learn more about plasma and the places you encounter it. Some samples of matter appear to have properties of solids, liquids, and/or gases at the same time. This can occur when the sample is composed of many small pieces. For example, we can pour sand as if it were a liquid because it is composed of many small grains of solid sand. Matter can also have properties of more than one state when it is a mixture, such as with clouds. Clouds appear to behave somewhat like gases, but they are actually mixtures of air (gas) and tiny particles of water (liquid or solid). The mass of an object is a measure of the amount of matter in it. One way to measure an object’s mass is to measure the force it takes to accelerate the object. It takes much more force to accelerate a car than a bicycle because the car has much more mass. A more common way to determine the mass of an object is to use a balance to compare its mass with a standard mass. Although weight is related to mass, it is not the same thing. Weight refers to the force that gravity exerts on an object. This force is directly proportional to the mass of the object. The weight of an object changes as the force of gravity changes, but its mass does not. An astronaut’s mass does not change just because she goes to the moon. But her weight on the moon is only one-sixth her earth-bound weight because the moon’s gravity is only one-sixth that of the earth’s. She may feel “weightless” during her trip when she experiences negligible external forces (gravitational or any other), although she is, of course, never “massless.” The law of conservation of matter summarizes many scientific observations about matter: It states that there is no detectable change in the total quantity of matter present when matter converts from one type to another (a chemical change) or changes among solid, liquid, or gaseous states (a physical change). Brewing beer and the operation of batteries provide examples of the conservation of matter (Figure 1.8). During the brewing of beer, the ingredients (water, yeast, grains, malt, hops, and sugar) are converted into beer (water, alcohol, carbonation, and flavoring substances) with no actual loss of substance. This is most clearly seen during the bottling process, when glucose turns This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 17 into ethanol and carbon dioxide, and the total mass of the substances does not change. This can also be seen in a lead-acid car battery: The original substances (lead, lead oxide, and sulfuric acid), which are capable of producing electricity, are changed into other substances (lead sulfate and water) that do not produce electricity, with no change in the actual amount of matter. Figure 1.8 (a) The mass of beer precursor materials is the same as the mass of beer produced: Sugar has become alcohol and carbonation. (b) The mass of the lead, lead oxide plates, and sulfuric acid that goes into the production of electricity is exactly equal to the mass of lead sulfate and water that is formed. Although this conservation law holds true for all conversions of matter, convincing examples are few and far between because, outside of the controlled conditions in a laboratory, we seldom collect all of the material that is produced during a particular conversion. For example, when you eat, digest, and assimilate food, all of the matter in the original food is preserved. But because some of the matter is incorporated into your body, and much is excreted as various types of waste, it is challenging to verify by measurement. Atoms and Molecules An atom is the smallest particle of an element that has the properties of that element and can enter into a chemical combination. Consider the element gold, for example. Imagine cutting a gold nugget in half, then cutting one of the halves in half, and repeating this process until a piece of gold remained that was so small that it could not be cut in half (regardless of how tiny your knife may be). This minimally sized piece of gold is an atom (from the Greek atomos, meaning “indivisible”) (Figure 1.9). This atom would no longer be gold if it were divided any further. 18 Chapter 1 | Essential Ideas Figure 1.9 (a) This photograph shows a gold nugget. (b) A scanning-tunneling microscope (STM) can generate views of the surfaces of solids, such as this image of a gold crystal. Each sphere represents one gold atom. (credit a: modification of work by United States Geological Survey; credit b: modification of work by “Erwinrossen”/Wikimedia Commons) The first suggestion that matter is composed of atoms is attributed to the Greek philosophers Leucippus and Democritus, who developed their ideas in the 5th century BCE. However, it was not until the early nineteenth century that John Dalton (1766–1844), a British schoolteacher with a keen interest in science, supported this hypothesis with quantitative measurements. Since that time, repeated experiments have confirmed many aspects of this hypothesis, and it has become one of the central theories of chemistry. Other aspects of Dalton’s atomic theory are still used but with minor revisions (details of Dalton’s theory are provided in the chapter on atoms and molecules). An atom is so small that its size is difficult to imagine. One of the smallest things we can see with our unaided eye is a single thread of a spider web: These strands are about 1/10,000 of a centimeter (0.00001 cm) in diameter. Although the cross-section of one strand is almost impossible to see without a microscope, it is huge on an atomic scale. A single carbon atom in the web has a diameter of about 0.000000015 centimeter, and it would take about 7000 carbon atoms to span the diameter of the strand. To put this in perspective, if a carbon atom were the size of a dime, the cross-section of one strand would be larger than a football field, which would require about 150 million carbon atom “dimes” to cover it. (Figure 1.10) shows increasingly close microscopic and atomic-level views of ordinary cotton. Figure 1.10 These images provide an increasingly closer view: (a) a cotton boll, (b) a single cotton fiber viewed under an optical microscope (magnified 40 times), (c) an image of a cotton fiber obtained with an electron microscope (much higher magnification than with the optical microscope); and (d and e) atomic-level models of the fiber (spheres of different colors represent atoms of different elements). (credit c: modification of work by “Featheredtar”/Wikimedia Commons) This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 19 An atom is so light that its mass is also difficult to imagine. A billion lead atoms (1,000,000,000 atoms) weigh about 3 × 10−13 grams, a mass that is far too light to be weighed on even the world’s most sensitive balances. It would require over 300,000,000,000,000 lead atoms (300 trillion, or 3 × 1014) to be weighed, and they would weigh only 0.0000001 gram. It is rare to find collections of individual atoms. Only a few elements, such as the gases helium, neon, and argon, consist of a collection of individual atoms that move about independently of one another. Other elements, such as the gases hydrogen, nitrogen, oxygen, and chlorine, are composed of units that consist of pairs of atoms (Figure 1.11). One form of the element phosphorus consists of units composed of four phosphorus atoms. The element sulfur exists in various forms, one of which consists of units composed of eight sulfur atoms. These units are called molecules. A molecule consists of two or more atoms joined by strong forces called chemical bonds. The atoms in a molecule move around as a unit, much like the cans of soda in a six-pack or a bunch of keys joined together on a single key ring. A molecule may consist of two or more identical atoms, as in the molecules found in the elements hydrogen, oxygen, and sulfur, or it may consist of two or more different atoms, as in the molecules found in water. Each water molecule is a unit that contains two hydrogen atoms and one oxygen atom. Each glucose molecule is a unit that contains 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. Like atoms, molecules are incredibly small and light. If an ordinary glass of water were enlarged to the size of the earth, the water molecules inside it would be about the size of golf balls. Figure 1.11 The elements hydrogen, oxygen, phosphorus, and sulfur form molecules consisting of two or more atoms of the same element. The compounds water, carbon dioxide, and glucose consist of combinations of atoms of different elements. Classifying Matter We can classify matter into several categories. Two broad categories are mixtures and pure substances. A pure substance has a constant composition. All specimens of a pure substance have exactly the same makeup and properties. Any sample of sucrose (table sugar) consists of 42.1% carbon, 6.5% hydrogen, and 51.4% oxygen by mass. Any sample of sucrose also has the same physical properties, such as melting point, color, and sweetness, regardless of the source from which it is isolated. We can divide pure substances into two classes: elements and compounds. Pure substances that cannot be broken down into simpler substances by chemical changes are called elements. Iron, silver, gold, aluminum, sulfur, oxygen, and copper are familiar examples of the more than 100 known elements, of which about 90 occur naturally on the earth, and two dozen or so have been created in laboratories. Pure substances that can be broken down by chemical changes are called compounds. This breakdown may produce either elements or other compounds, or both. Mercury(II) oxide, an orange, crystalline solid, can be broken down by heat into the elements mercury and oxygen (Figure 1.12). When heated in the absence of air, the compound sucrose is broken down into the element carbon and the compound water. (The initial stage of this process, when the sugar is 20 Chapter 1 | Essential Ideas turning brown, is known as caramelization—this is what imparts the characteristic sweet and nutty flavor to caramel apples, caramelized onions, and caramel). Silver(I) chloride is a white solid that can be broken down into its elements, silver and chlorine, by absorption of light. This property is the basis for the use of this compound in photographic films and photochromic eyeglasses (those with lenses that darken when exposed to light). Figure 1.12 (a)The compound mercury(II) oxide, (b)when heated, (c) decomposes into silvery droplets of liquid mercury and invisible oxygen gas. (credit: modification of work by Paul Flowers) Link to Learning Many compounds break down when heated. This site (http://openstaxcollege.org/l/16mercury) shows the breakdown of mercury oxide, HgO. You can also view an example of the photochemical decomposition of silver chloride (http://openstaxcollege.org/l/16silvchloride) (AgCl), the basis of early photography. The properties of combined elements are different from those in the free, or uncombined, state. For example, white crystalline sugar (sucrose) is a compound resulting from the chemical combination of the element carbon, which is a black solid in one of its uncombined forms, and the two elements hydrogen and oxygen, which are colorless gases when uncombined. Free sodium, an element that is a soft, shiny, metallic solid, and free chlorine, an element that is a yellow-green gas, combine to form sodium chloride (table salt), a compound that is a white, crystalline solid. A mixture is composed of two or more types of matter that can be present in varying amounts and can be separated by physical changes, such as evaporation (you will learn more about this later). A mixture with a composition that varies from point to point is called a heterogeneous mixture. Italian dressing is an example of a heterogeneous mixture (Figure 1.13). Its composition can vary because we can make it from varying amounts of oil, vinegar, and herbs. It is not the same from point to point throughout the mixture—one drop may be mostly vinegar, whereas a different drop may be mostly oil or herbs because the oil and vinegar separate and the herbs settle. Other examples of heterogeneous mixtures are chocolate chip cookies (we can see the separate bits of chocolate, nuts, and cookie dough) and granite (we can see the quartz, mica, feldspar, and more). A homogeneous mixture, also called a solution, exhibits a uniform composition and appears visually the same throughout. An example of a solution is a sports drink, consisting of water, sugar, coloring, flavoring, and electrolytes mixed together uniformly (Figure 1.13). Each drop of a sports drink tastes the same because each drop contains the same amounts of water, sugar, and other components. Note that the composition of a sports drink can vary—it could be made with somewhat more or less sugar, flavoring, or other components, and still be a sports drink. Other examples of homogeneous mixtures include air, maple syrup, gasoline, and a solution of salt in water. This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 21 Figure 1.13 (a) Oil and vinegar salad dressing is a heterogeneous mixture because its composition is not uniform throughout. (b) A commercial sports drink is a homogeneous mixture because its composition is uniform throughout. (credit a “left”: modification of work by John Mayer; credit a “right”: modification of work by Umberto Salvagnin; credit b “left: modification of work by Jeff Bedford) Although there are just over 100 elements, tens of millions of chemical compounds result from different combinations of these elements. Each compound has a specific composition and possesses definite chemical and physical properties by which we can distinguish it from all other compounds. And, of course, there are innumerable ways to combine elements and compounds to form different mixtures. A summary of how to distinguish between the various major classifications of matter is shown in (Figure 1.14). Figure 1.14 Depending on its properties, a given substance can be classified as a homogeneous mixture, a heterogeneous mixture, a compound, or an element. Eleven elements make up about 99% of the earth’s crust and atmosphere (Table 1.1). Oxygen constitutes nearly one- half and silicon about one-quarter of the total quantity of these elements. A majority of elements on earth are found in chemical combinations with other elements; about one-quarter of the elements are also found in the free state. Elemental Composition of Earth Element Symbol Percent Mass Element Symbol Percent Mass oxygen O 49.20 chlorine Cl 0.19 silicon Si 25.67 phosphorus P 0.11 Table 1.1 22 Chapter 1 | Essential Ideas Elemental Composition of Earth Element Symbol Percent Mass Element Symbol Percent Mass aluminum Al 7.50 manganese Mn 0.09 iron Fe 4.71 carbon C 0.08 calcium Ca 3.39 sulfur S 0.06 sodium Na 2.63 barium Ba 0.04 potassium K 2.40 nitrogen N 0.03 magnesium Mg 1.93 fluorine F 0.03 hydrogen H 0.87 strontium Sr 0.02 titanium Ti 0.58 all others - 0.47 Table 1.1 Chemistry in Everyday Life Decomposition of Water / Production of Hydrogen Water consists of the elements hydrogen and oxygen combined in a 2 to 1 ratio. Water can be broken down into hydrogen and oxygen gases by the addition of energy. One way to do this is with a battery or power supply, as shown in (Figure 1.15). This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 23 Figure 1.15 The decomposition of water is shown at the macroscopic, microscopic, and symbolic levels. The battery provides an electric current (microscopic) that decomposes water. At the macroscopic level, the liquid separates into the gases hydrogen (on the left) and oxygen (on the right). Symbolically, this change is presented by showing how liquid H2O separates into H2 and O2 gases. The breakdown of water involves a rearrangement of the atoms in water molecules into different molecules, each composed of two hydrogen atoms and two oxygen atoms, respectively. Two water molecules form one oxygen molecule and two hydrogen molecules. The representation for what occurs, 2H 2 O(l) ⟶ 2H 2(g) + O 2(g), will be explored in more depth in later chapters. The two gases produced have distinctly different properties. Oxygen is not flammable but is required for combustion of a fuel, and hydrogen is highly flammable and a potent energy source. How might this knowledge be applied in our world? One application involves research into more fuel-efficient transportation. Fuel-cell vehicles (FCV) run on hydrogen instead of gasoline (Figure 1.16). They are more efficient than vehicles with internal combustion engines, are nonpolluting, and reduce greenhouse gas emissions, making us less dependent on fossil fuels. FCVs are not yet economically viable, however, and current hydrogen production depends on natural gas. If we can develop a process to economically decompose water, or produce hydrogen in another environmentally sound way, FCVs may be the way of the future. 24 Chapter 1 | Essential Ideas Figure 1.16 A fuel cell generates electrical energy from hydrogen and oxygen via an electrochemical process and produces only water as the waste product. Chemistry in Everyday Life Chemistry of Cell Phones Imagine how different your life would be without cell phones (Figure 1.17) and other smart devices. Cell phones are made from numerous chemical substances, which are extracted, refined, purified, and assembled using an extensive and in-depth understanding of chemical principles. About 30% of the elements that are found in nature are found within a typical smart phone. The case/body/frame consists of a combination of sturdy, durable polymers comprised primarily of carbon, hydrogen, oxygen, and nitrogen [acrylonitrile butadiene styrene (ABS) and polycarbonate thermoplastics], and light, strong, structural metals, such as aluminum, magnesium, and iron. The display screen is made from a specially toughened glass (silica glass strengthened by the addition of aluminum, sodium, and potassium) and coated with a material to make it conductive (such as indium tin oxide). The circuit board uses a semiconductor material (usually silicon); commonly used metals like copper, tin, silver, and gold; and more unfamiliar elements such as yttrium, praseodymium, and gadolinium. The battery relies upon lithium ions and a variety of other materials, including iron, cobalt, copper, polyethylene oxide, and polyacrylonitrile. This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 25 Figure 1.17 Almost one-third of naturally occurring elements are used to make a cell phone. (credit: modification of work by John Taylor) 1.3 Physical and Chemical Properties By the end of this section, you will be able to: Identify properties of and changes in matter as physical or chemical Identify properties of matter as extensive or intensive The characteristics that enable us to distinguish one substance from another are called properties. A physical property is a characteristic of matter that is not associated with a change in its chemical composition. Familiar examples of physical properties include density, color, hardness, melting and boiling points, and electrical conductivity. We can observe some physical properties, such as density and color, without changing the physical state of the matter observed. Other physical properties, such as the melting temperature of iron or the freezing temperature of water, can only be observed as matter undergoes a physical change. A physical change is a change in the state or properties of matter without any accompanying change in its chemical composition (the identities of the substances contained in the matter). We observe a physical change when wax melts, when sugar dissolves in coffee, and when steam condenses into liquid water (Figure 1.18). Other examples of physical changes include magnetizing and demagnetizing metals (as is done with common antitheft security tags) and grinding solids into powders (which can sometimes yield noticeable changes in color). In each of these examples, there is a change in the physical state, form, or properties of the substance, but no change in its chemical composition. 26 Chapter 1 | Essential Ideas Figure 1.18 (a) Wax undergoes a physical change when solid wax is heated and forms liquid wax. (b) Steam condensing inside a cooking pot is a physical change, as water vapor is changed into liquid water. (credit a: modification of work by “95jb14”/Wikimedia Commons; credit b: modification of work by “mjneuby”/Flickr) The change of one type of matter into another type (or the inability to change) is a chemical property. Examples of chemical properties include flammability, toxicity, acidity, reactivity (many types), and heat of combustion. Iron, for example, combines with oxygen in the presence of water to form rust; chromium does not oxidize (Figure 1.19). Nitroglycerin is very dangerous because it explodes easily; neon poses almost no hazard because it is very unreactive. Figure 1.19 (a) One of the chemical properties of iron is that it rusts; (b) one of the chemical properties of chromium is that it does not. (credit a: modification of work by Tony Hisgett; credit b: modification of work by “Atoma”/Wikimedia Commons) To identify a chemical property, we look for a chemical change. A chemical change always produces one or more types of matter that differ from the matter present before the change. The formation of rust is a chemical change because rust is a different kind of matter than the iron, oxygen, and water present before the rust formed. The explosion of nitroglycerin is a chemical change because the gases produced are very different kinds of matter from the original substance. Other examples of chemical changes include reactions that are performed in a lab (such as copper reacting with nitric acid), all forms of combustion (burning), and food being cooked, digested, or rotting (Figure 1.20). This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 27 Figure 1.20 (a) Copper and nitric acid undergo a chemical change to form copper nitrate and brown, gaseous nitrogen dioxide. (b) During the combustion of a match, cellulose in the match and oxygen from the air undergo a chemical change to form carbon dioxide and water vapor. (c) Cooking red meat causes a number of chemical changes, including the oxidation of iron in myoglobin that results in the familiar red-to-brown color change. (d) A banana turning brown is a chemical change as new, darker (and less tasty) substances form. (credit b: modification of work by Jeff Turner; credit c: modification of work by Gloria Cabada-Leman; credit d: modification of work by Roberto Verzo) Properties of matter fall into one of two categories. If the property depends on the amount of matter present, it is an extensive property. The mass and volume of a substance are examples of extensive properties; for instance, a gallon of milk has a larger mass and volume than a cup of milk. The value of an extensive property is directly proportional to the amount of matter in question. If the property of a sample of matter does not depend on the amount of matter present, it is an intensive property. Temperature is an example of an intensive property. If the gallon and cup of milk are each at 20 °C (room temperature), when they are combined, the temperature remains at 20 °C. As another example, consider the distinct but related properties of heat and temperature. A drop of hot cooking oil spattered on your arm causes brief, minor discomfort, whereas a pot of hot oil yields severe burns. Both the drop and the pot of oil are at the same temperature (an intensive property), but the pot clearly contains much more heat (extensive property). Chemistry in Everyday Life Hazard Diamond You may have seen the symbol shown in Figure 1.21 on containers of chemicals in a laboratory or workplace. Sometimes called a “fire diamond” or “hazard diamond,” this chemical hazard diamond provides valuable 28 Chapter 1 | Essential Ideas information that briefly summarizes the various dangers of which to be aware when working with a particular substance. Figure 1.21 The National Fire Protection Agency (NFPA) hazard diamond summarizes the major hazards of a chemical substance. The National Fire Protection Agency (NFPA) 704 Hazard Identification System was developed by NFPA to provide safety information about certain substances. The system details flammability, reactivity, health, and other hazards. Within the overall diamond symbol, the top (red) diamond specifies the level of fire hazard (temperature range for flash point). The blue (left) diamond indicates the level of health hazard. The yellow (right) diamond describes reactivity hazards, such as how readily the substance will undergo detonation or a violent chemical change. The white (bottom) diamond points out special hazards, such as if it is an oxidizer (which allows the substance to burn in the absence of air/oxygen), undergoes an unusual or dangerous reaction with water, is corrosive, acidic, alkaline, a biological hazard, radioactive, and so on. Each hazard is rated on a scale from 0 to 4, with 0 being no hazard and 4 being extremely hazardous. While many elements differ dramatically in their chemical and physical properties, some elements have similar properties. We can identify sets of elements that exhibit common behaviors. For example, many elements conduct heat and electricity well, whereas others are poor conductors. These properties can be used to sort the elements into three classes: metals (elements that conduct well), nonmetals (elements that conduct poorly), and metalloids (elements that have properties of both metals and nonmetals). The periodic table is a table of elements that places elements with similar properties close together (Figure 1.22). You will learn more about the periodic table as you continue your study of chemistry. This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 29 Figure 1.22 The periodic table shows how elements may be grouped according to certain similar properties. Note the background color denotes whether an element is a metal, metalloid, or nonmetal, whereas the element symbol color indicates whether it is a solid, liquid, or gas. 1.4 Measurements By the end of this section, you will be able to: Explain the process of measurement Identify the three basic parts of a quantity Describe the properties and units of length, mass, volume, density, temperature, and time Perform basic unit calculations and conversions in the metric and other unit systems Measurements provide the macroscopic information that is the basis of most of the hypotheses, theories, and laws that describe the behavior of matter and energy in both the macroscopic and microscopic domains of chemistry. Every measurement provides three kinds of information: the size or magnitude of the measurement (a number); a standard of comparison for the measurement (a unit); and an indication of the uncertainty of the measurement. While the number and unit are explicitly represented when a quantity is written, the uncertainty is an aspect of the measurement result that is more implicitly represented and will be discussed later. 30 Chapter 1 | Essential Ideas The number in the measurement can be represented in different ways, including decimal form and scientific notation. (Scientific notation is also known as exponential notation; a review of this topic can be found in Appendix B.) For example, the maximum takeoff weight of a Boeing 777-200ER airliner is 298,000 kilograms, which can also be written as 2.98 × 105 kg. The mass of the average mosquito is about 0.0000025 kilograms, which can be written as 2.5 × 10−6 kg. Units, such as liters, pounds, and centimeters, are standards of comparison for measurements. When we buy a 2-liter bottle of a soft drink, we expect that the volume of the drink was measured, so it is two times larger than the volume that everyone agrees to be 1 liter. The meat used to prepare a 0.25-pound hamburger is measured so it weighs one- fourth as much as 1 pound. Without units, a number can be meaningless, confusing, or possibly life threatening. Suppose a doctor prescribes phenobarbital to control a patient’s seizures and states a dosage of “100” without specifying units. Not only will this be confusing to the medical professional giving the dose, but the consequences can be dire: 100 mg given three times per day can be effective as an anticonvulsant, but a single dose of 100 g is more than 10 times the lethal amount. We usually report the results of scientific measurements in SI units, an updated version of the metric system, using the units listed in Table 1.2. Other units can be derived from these base units. The standards for these units are fixed by international agreement, and they are called the International System of Units or SI Units (from the French, Le Système International d’Unités). SI units have been used by the United States National Institute of Standards and Technology (NIST) since 1964. Base Units of the SI System Property Measured Name of Unit Symbol of Unit length meter m mass kilogram kg time second s temperature kelvin K electric current ampere A amount of substance mole mol luminous intensity candela cd Table 1.2 Sometimes we use units that are fractions or multiples of a base unit. Ice cream is sold in quarts (a familiar, non-SI base unit), pints (0.5 quart), or gallons (4 quarts). We also use fractions or multiples of units in the SI system, but these fractions or multiples are always powers of 10. Fractional or multiple SI units are named using a prefix and the name of the base unit. For example, a length of 1000 meters is also called a kilometer because the prefix kilo means “one thousand,” which in scientific notation is 103 (1 kilometer = 1000 m = 103 m). The prefixes used and the powers to which 10 are raised are listed in Table 1.3. This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 31 Common Unit Prefixes Prefix Symbol Factor Example femto f 10−15 1 femtosecond (fs) = 1 × 10−15 m (0.000000000000001 s) pico p 10−12 1 picometer (pm) = 1 × 10−12 m (0.000000000001 m) nano n 10−9 4 nanograms (ng) = 4 × 10−9 g (0.000000004 g) micro µ 10−6 1 microliter (μL) = 1 × 10−6 L (0.000001 L) milli m 10−3 2 millimoles (mmol) = 2 × 10−3 mol (0.002 mol) centi c 10−2 7 centimeters (cm) = 7 × 10−2 m (0.07 m) deci d 10−1 1 deciliter (dL) = 1 × 10−1 L (0.1 L ) kilo k 103 1 kilometer (km) = 1 × 103 m (1000 m) mega M 106 3 megahertz (MHz) = 3 × 106 Hz (3,000,000 Hz) giga G 109 8 gigayears (Gyr) = 8 × 109 yr (8,000,000,000 Gyr) tera T 1012 5 terawatts (TW) = 5 × 1012 W (5,000,000,000,000 W) Table 1.3 Link to Learning Need a refresher or more practice with scientific notation? Visit this site (http://openstaxcollege.org/l/16notation) to go over the basics of scientific notation. SI Base Units The initial units of the metric system, which eventually evolved into the SI system, were established in France during the French Revolution. The original standards for the meter and the kilogram were adopted there in 1799 and eventually by other countries. This section introduces four of the SI base units commonly used in chemistry. Other SI units will be introduced in subsequent chapters. Length The standard unit of length in both the SI and original metric systems is the meter (m). A meter was originally specified as 1/10,000,000 of the distance from the North Pole to the equator. It is now defined as the distance light in a vacuum travels in 1/299,792,458 of a second. A meter is about 3 inches longer than a yard (Figure 1.23); one 32 Chapter 1 | Essential Ideas meter is about 39.37 inches or 1.094 yards. Longer distances are often reported in kilometers (1 km = 1000 m = 103 m), whereas shorter distances can be reported in centimeters (1 cm = 0.01 m = 10−2 m) or millimeters (1 mm = 0.001 m = 10−3 m). Figure 1.23 The relative lengths of 1 m, 1 yd, 1 cm, and 1 in. are shown (not actual size), as well as comparisons of 2.54 cm and 1 in., and of 1 m and 1.094 yd. Mass The standard unit of mass in the SI system is the kilogram (kg). A kilogram was originally defined as the mass of a liter of water (a cube of water with an edge length of exactly 0.1 meter). It is now defined by a certain cylinder of platinum-iridium alloy, which is kept in France (Figure 1.24). Any object with the same mass as this cylinder is said to have a mass of 1 kilogram. One kilogram is about 2.2 pounds. The gram (g) is exactly equal to 1/1000 of the mass of the kilogram (10−3 kg). Figure 1.24 This replica prototype kilogram is housed at the National Institute of Standards and Technology (NIST) in Maryland. (credit: National Institutes of Standards and Technology) This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 33 Temperature Temperature is an intensive property. The SI unit of temperature is the kelvin (K). The IUPAC convention is to use kelvin (all lowercase) for the word, K (uppercase) for the unit symbol, and neither the word “degree” nor the degree symbol (°). The degree Celsius (°C) is also allowed in the SI system, with both the word “degree” and the degree symbol used for Celsius measurements. Celsius degrees are the same magnitude as those of kelvin, but the two scales place their zeros in different places. Water freezes at 273.15 K (0 °C) and boils at 373.15 K (100 °C) by definition, and normal human body temperature is approximately 310 K (37 °C). The conversion between these two units and the Fahrenheit scale will be discussed later in this chapter. Time The SI base unit of time is the second (s). Small and large time intervals can be expressed with the appropriate prefixes; for example, 3 microseconds = 0.000003 s = 3 × 10−6 and 5 megaseconds = 5,000,000 s = 5 × 106 s. Alternatively, hours, days, and years can be used. Derived SI Units We can derive many units from the seven SI base units. For example, we can use the base unit of length to define a unit of volume, and the base units of mass and length to define a unit of density. Volume Volume is the measure of the amount of space occupied by an object. The standard SI unit of volume is defined by the base unit of length (Figure 1.25). The standard volume is a cubic meter (m3), a cube with an edge length of exactly one meter. To dispense a cubic meter of water, we could build a cubic box with edge lengths of exactly one meter. This box would hold a cubic meter of water or any other substance. A more commonly used unit of volume is derived from the decimeter (0.1 m, or 10 cm). A cube with edge lengths of exactly one decimeter contains a volume of one cubic decimeter (dm3). A liter (L) is the more common name for the cubic decimeter. One liter is about 1.06 quarts. A cubic centimeter (cm3) is the volume of a cube with an edge length of exactly one centimeter. The abbreviation cc (for cubic centimeter) is often used by health professionals. A cubic centimeter is also called a milliliter (mL) and is 1/1000 of a liter. 34 Chapter 1 | Essential Ideas Figure 1.25 (a) The relative volumes are shown for cubes of 1 m3, 1 dm3 (1 L), and 1 cm3 (1 mL) (not to scale). (b) The diameter of a dime is compared relative to the edge length of a 1-cm3 (1-mL) cube. Density We use the mass and volume of a substance to determine its density. Thus, the units of density are defined by the base units of mass and length. The density of a substance is the ratio of the mass of a sample of the substance to its volume. The SI unit for density is the kilogram per cubic meter (kg/m3). For many situations, however, this as an inconvenient unit, and we often use grams per cubic centimeter (g/cm3) for the densities of solids and liquids, and grams per liter (g/L) for gases. Although there are exceptions, most liquids and solids have densities that range from about 0.7 g/cm3 (the density of gasoline) to 19 g/cm3 (the density of gold). The density of air is about 1.2 g/L. Table 1.4 shows the densities of some common substances. Densities of Common Substances Solids Liquids Gases (at 25 °C and 1 atm) ice (at 0 °C) 0.92 g/cm3 water 1.0 g/cm3 dry air 1.20 g/L oak (wood) 0.60–0.90 g/cm3 ethanol 0.79 g/cm3 oxygen 1.31 g/L iron 7.9 g/cm3 acetone 0.79 g/cm3 nitrogen 1.14 g/L copper 9.0 g/cm3 glycerin 1.26 g/cm3 carbon dioxide 1.80 g/L lead 11.3 g/cm3 olive oil 0.92 g/cm3 helium 0.16 g/L Table 1.4 This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 35 Densities of Common Substances Solids Liquids Gases (at 25 °C and 1 atm) silver 10.5 g/cm3 gasoline 0.70–0.77 g/cm3 neon 0.83 g/L gold 19.3 g/cm3 mercury 13.6 g/cm3 radon 9.1 g/L Table 1.4 While there are many ways to determine the density of an object, perhaps the most straightforward method involves separately finding the mass and volume of the object, and then dividing the mass of the sample by its volume. In the following example, the mass is found directly by weighing, but the volume is found indirectly through length measurements. density = mass volume Example 1.1 Calculation of Density Gold—in bricks, bars, and coins—has been a form of currency for centuries. In order to swindle people into paying for a brick of gold without actually investing in a brick of gold, people have considered filling the centers of hollow gold bricks with lead to fool buyers into thinking that the entire brick is gold. It does not work: Lead is a dense substance, but its density is not as great as that of gold, 19.3 g/cm3. What is the density of lead if a cube of lead has an edge length of 2.00 cm and a mass of 90.7 g? Solution The density of a substance can be calculated by dividing its mass by its volume. The volume of a cube is calculated by cubing the edge length. volume of lead cube = 2.00 cm × 2.00 cm × 2.00 cm = 8.00 cm 3 90.7 g 11.3 g density = mass = = = 11.3 g/cm 3 volume 8.00 cm 3 1.00 cm 3 (We will discuss the reason for rounding to the first decimal place in the next section.) Check Your Learning (a) To three decimal places, what is the volume of a cube (cm3) with an edge length of 0.843 cm? (b) If the cube in part (a) is copper and has a mass of 5.34 g, what is the density of copper to two decimal places? Answer: (a) 0.599 cm3; (b) 8.91 g/cm3 36 Chapter 1 | Essential Ideas Link to Learning To learn more about the relationship between mass, volume, and density, use this interactive simulator (http://openstaxcollege.org/l/16phetmasvolden) to explore the density of different materials, like wood, ice, brick, and aluminum. Example 1.2 Using Displacement of Water to Determine Density This PhET simulation (http://openstaxcollege.org/l/16phetmasvolden) illustrates another way to determine density, using displacement of water. Determine the density of the red and yellow blocks. Solution When you open the density simulation and select Same Mass, you can choose from several 5.00-kg colored blocks that you can drop into a tank containing 100.00 L water. The yellow block floats (it is less dense than water), and the water level rises to 105.00 L. While floating, the yellow block displaces 5.00 L water, an amount equal to the weight of the block. The red block sinks (it is more dense than water, which has density = 1.00 kg/L), and the water level rises to 101.25 L. The red block therefore displaces 1.25 L water, an amount equal to the volume of the block. The density of the red block is: density = mass = 5.00 kg = 4.00 kg/L volume 1.25 L Note that since the yellow block is not completely submerged, you cannot determine its density from this information. But if you hold the yellow block on the bottom of the tank, the water level rises to 110.00 L, which means that it now displaces 10.00 L water, and its density can be found: density = mass = 5.00 kg = 0.500 kg/L volume 10.00 L Check Your Learning Remove all of the blocks from the water and add the green block to the tank of water, placing it approximately in the middle of the tank. Determine the density of the green block. Answer: 2.00 kg/L 1.5 Measurement Uncertainty, Accuracy, and Precision By the end of this section, you will be able to: Define accuracy and precision Distinguish exact and uncertain numbers Correctly represent uncertainty in quantities using significant figures Apply proper rounding rules to computed quantities Counting is the only type of measurement that is free from uncertainty, provided the number of objects being counted does not change while the counting process is underway. The result of such a counting measurement is an example of an exact number. If we count eggs in a carton, we know exactly how many eggs the carton contains. The numbers of defined quantities are also exact. By definition, 1 foot is exactly 12 inches, 1 inch is exactly 2.54 centimeters, and This content is available for free at https://cnx.org/content/col11760/1.9 Chapter 1 | Essential Ideas 37 1 gram is exactly 0.001 kilogram. Quantities derived from measurements other t

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