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SCHAUM’S OUTLINE OF Theory and Problems of BEGINNING CHEMISTRY Third Edition David E. Goldberg, Ph.D. Professor of Chemistry Brooklyn College City University of New York Schaum’s Outline Series McGRAW-HILL New York Chicago San Francisco Lisbon London Madrid Mexico City Milan New Delhi San Juan Seoul Singapore Sydney Toronto Copyright © 2005, 1999, 1991 by The McGraw-Hill Companies, Inc. All rights reserved. Manufactured in the United States of America. Except as permitted under the United States Copyright Act of 1976, no part of this publication may be reproduced or distributed in any form or by any means, or stored in a database or retrieval system, without the prior written permission of the publisher. 0-07-146628-2 The material in this eBook also appears in the print version of this title: 0-07-144780-6. All trademarks are trademarks of their respective owners. Rather than put a trademark symbol after every occurrence of a trademarked name, we use names in an editorial fashion only, and to the benefit of the trademark owner, with no intention of infringement of the trademark. Where such designations appear in this book, they have been printed with initial caps. McGraw-Hill eBooks are available at special quantity discounts to use as premiums and sales promotions, or for use in corporate training programs. For more information, please contact George Hoare, Special Sales, at george_hoaremcgraw-hill.com or (212) 904-4069. TERMS OF USE This is a copyrighted work and The McGraw-Hill Companies, Inc. (“McGraw-Hill”) and its licensors reserve all rights in and to the work. Use of this work is subject to these terms. Except as permitted under the Copyright Act of 1976 and the right to store and retrieve one copy of the work, you may not decompile, disassemble, reverse engineer, reproduce, modify, create derivative works based upon, transmit, distribute, disseminate, sell, publish or sublicense the work or any part of it without McGraw- Hill’s prior consent. You may use the work for your own noncommercial and personal use; any other use of the work is strictly prohibited. Your right to use the work may be terminated if you fail to comply with these terms. THE WORK IS PROVIDED “AS IS.” McGRAW-HILL AND ITS LICENSORS MAKE NO GUARANTEES OR WARRANTIES AS TO THE ACCURACY, ADEQUACY OR COMPLETENESS OF OR RESULTS TO BE OBTAINED FROM USING THE WORK, INCLUDING ANY INFORMATION THAT CAN BE ACCESSED THROUGH THE WORK VIA HYPERLINK OR OTHERWISE, AND EXPRESSLY DISCLAIM ANY WARRANTY, EXPRESS OR IMPLIED, INCLUDING BUT NOT LIMITED TO IMPLIED WARRANTIES OF MERCHANTABILITY OR FITNESS FOR A PARTICULAR PURPOSE. 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DOI: 10.1036/0071466282 PREFACE This book is designed to help students do well in their first chemistry course, especially those who have little or no chemistry background. It can be used effectively in a course preparatory to a general college chemistry course as well as in a course in chemistry for liberal arts students. It should also provide additional assistance to students in the first semester of a chemistry course for nurses and others in the allied health fields. It will prove to be of value in a high school chemistry course and in a general chemistry course for majors. The book aims to help the student develop both problem-solving skills and skill in precise reading and interpreting scientific problems and questions. Analogies to everyday life introduce certain types of problems to make the underlying principles less abstract. Many of the problems were devised to clarify particular points often confused by beginning students. To ensure mastery, the book often presents problems in parts, then asks the same question as an entity, to see if the student can do the parts without the aid of the fragmented question. It provides some figures that have proved helpful to a generation of students. The author gratefully acknowledges the help of the editors at McGraw-Hill. DAVID E. GOLDBERG iii Copyright © 2005, 1999, 1991 by The McGraw-Hill Companies, Inc. Click here for terms of use. This page is intentionally left blank. TO THE STUDENT This book is designed to help you understand chemistry fundamentals. Learning chemistry requires that you master chemical terminology and be able to perform calculations with ease. Toward these ends, many of the examples and problems are formulated to alert you to questions that sound different but are actually the same (Problem 3.16 for example) or questions that are different but sound very similar (Problems 5.13 and 7.25, for example). You should not attempt to memorize the solutions to the problems. (There is enough to memorize, without that.) Instead, you must try to understand the concepts involved. Your instructor and texts usually teach generalities (e.g., Atoms of all main group elements except noble gases have the number of outermost electrons equal to their group number.), but the instructor asks specific questions on exams (e.g., How many outermost electrons are there in a phosphorus atom?) You must not only know the principle, but also in what situations it applies. You must practice by working many problems, because in addition to the principles, you must get accustomed to the many details involved in solving problems correctly. The key to success in chemistry is working very many problems To get the most from this book, use a 5 × 8 card to cover up the solutions while you are doing the problems. Do not look at the answer first. It is easy to convince yourself that you know how to do a problem by looking at the answer, but generating the answer yourself, as you must do on exams, is not the same. After you have finished, compare your result with the answer given. If the method differs, it does not mean that your method is necessarily incorrect. If your answer is the same, your method is probably correct. Otherwise, try to understand what the difference is, and where you made a mistake, if you did so. Some of the problems given after the text are very short and/or very easy (Problems 5.12 and 5.14, for example). They are designed to emphasize a particular point. After you get the correct answer, ask yourself why such a question was asked. Many other problems give analogies to everyday life, to help you understand a chemical principle (Problems 2.13 with 2.14, 4.6, 5.15 with 5.16, 7.13 through 7.16 and 10.41, for example). Make sure you understand the chemical meaning of the terms presented throughout the semester. For example, “significant figures” means something very different in chemical calculations than in economic discussions. Special terms used for the first time in this book will be italicized. Whenever you encounter such a term, use it repeatedly until you thoroughly understand its meaning. If necessary, use the Glossary to find the meanings of unfamiliar terms. Always use the proper units with measurable quantities. It makes quite a bit of difference if your pet is 4 in. tall or 4 ft tall After Chapter 2, always use the proper number of significant figures in your calculations. Do yourself a favor and use the same symbols and abbreviations for chemical quantities that are used in the text. If you use a different symbol, you might become confused later when that symbol is used for a different quantity. Some of the problems are stated in parts. After you do the problem by solving the various parts, see if you would know how to solve the same problem if only the last part were asked. The conversion figure on page 348 shows all the conversions presented in the book. As you proceed, add the current conversions from the figure to your solution techniques. v Copyright © 2005, 1999, 1991 by The McGraw-Hill Companies, Inc. Click here for terms of use. This page is intentionally left blank. For more information about this title, click here CONTENTS CHAPTER 1 Basic Concepts 1 1.1 Introduction 1 1.2 The Elements 1 1.3 Matter and Energy 2 1.4 Properties 3 1.5 Classification of Matter 3 1.6 Representation of Elements 5 1.7 Laws, Hypotheses, and Theories 6 CHAPTER 2 Mathematical Methods in Chemistry 10 2.1 Introduction 10 2.2 Factor-Label Method 10 2.3 Metric System 12 2.4 Exponential Numbers 16 2.5 Significant Digits 17 2.6 Density 21 2.7 Temperature Scales 23 CHAPTER 3 Atoms and Atomic Masses 38 3.1 Introduction 38 3.2 Atomic Theory 38 3.3 Atomic Masses 39 3.4 Atomic Structure 40 3.5 Isotopes 41 3.6 Periodic Table 42 Chapter 4 Electronic Configuration of the Atom 51 4.1 Introduction 51 4.2 Bohr Theory 51 4.3 Quantum Numbers 53 4.4 Quantum Numbers and Energies of Electrons 54 4.5 Shells, Subshells, and Orbitals 55 4.6 Shapes of Orbitals 58 4.7 Buildup Principle 58 4.8 Electronic Structure and the Periodic Table 60 viiviii CONTENTS Chapter 5 Chemical Bonding 67 5.1 Introduction 67 5.2 Chemical Formulas 67 5.3 The Octet Rule 68 5.4 Ions 69 5.5 Electron Dot Notation 71 5.6 Covalent Bonding 72 5.7 Distinction Between Ionic and Covalent Bonding 74 5.8 Predicting the Nature of Bonding in Compounds 75 5.9 Detailed Electronic Configurations of Ions (Optional) 76 Chapter 6 Inorganic Nomenclature 86 6.1 Introduction 86 6.2 Binary Compounds of Nonmetals 87 6.3 Naming Ionic Compounds 88 6.4 Naming Inorganic Acids 93 6.5 Acid Salts 94 6.6 Hydrates 94 Chapter 7 Formula Calculations 102 7.1 Introduction 102 7.2 Molecules and Formula Units 102 7.3 Formula Masses 103 7.4 The Mole 103 7.5 Percent Composition of Compounds 106 7.6 Empirical Formulas 107 7.7 Molecular Formulas 108 Chapter 8 Chemical Equations 120 8.1 Introduction 120 8.2 Balancing Simple Equations 121 8.3 Predicting the Products of a Reaction 122 Chapter 9 Net Ionic Equations 134 9.1 Introduction 134 9.2 Writing Net Ionic Equations 134 Chapter 10 Stoichiometry 142 10.1 Mole-to-Mole Calculations 142 10.2 Calculations Involving Other Quantities 143 10.3 Limiting Quantities 144 10.4 Calculations Based on Net Ionic Equations 147 10.5 Heat Capacity and Heat of Reaction 147CONTENTS ix Chapter 11 Molarity 162 11.1 Introduction 162 11.2 Molarity Calculations 162 11.3 Titration 164 11.4 Stoichiometry in Solution 166 Chapter 12 Gases 173 12.1 Introduction 173 12.2 Pressure of Gases 173 12.3 Boyle’s Law 174 12.4 Graphical Representation of Data 175 12.5 Charles’ Law 177 12.6 The Combined Gas Law 180 12.7 The Ideal Gas Law 181 12.8 Dalton’s Law of Partial Pressures 183 Chapter 13 Kinetic Molecular Theory 195 13.1 Introduction 195 13.2 Postulates of the Kinetic Molecular Theory 195 13.3 Explanation of Gas Pressure, Boyle’s Law, and Charles’ Law 196 13.4 Graham’s Law 197 Chapter 14 Oxidation and Reduction 201 14.1 Introduction 201 14.2 Assigning Oxidation Numbers 201 14.3 Periodic Relationships of Oxidation Numbers 203 14.4 Oxidation Numbers in Inorganic Nomenclature 205 14.5 Balancing Oxidation-Reduction Equations 205 14.6 Electrochemistry 209 Chapter 15 Solutions 219 15.1 Qualitative Concentration Terms 219 15.2 Molality 219 15.3 Mole Fraction 220 15.4 Equivalents 221 15.5 Normality 222 15.6 Equivalent Mass 223 Chapter 16 Rates and Equilibrium 230 16.1 Introduction 230 16.2 Rates of Chemical Reaction 230 16.3 Chemical Equilibrium 232 16.4 Equilibrium Constants 234x CONTENTS Chapter 17 Acid-Base Theory 246 17.1 Introduction 246 17.2 The Brønsted-Lowry Theory 246 17.3 Acid-Base Equilibrium 248 17.4 Autoionization of Water 249 17.5 The pH Scale 250 17.6 Buffer Solutions 251 Chapter 18 Organic Chemistry 261 18.1 Introduction 261 18.2 Bonding in Organic Compounds 261 18.3 Structural, Condensed, and Line Formulas 262 18.4 Hydrocarbons 264 18.5 Isomerism 266 18.6 Radicals and Functional Groups 267 18.7 Alcohols 269 18.8 Ethers 270 18.9 Aldehydes and Ketones 271 18.10 Acids and Esters 271 18.11 Amines 272 18.12 Amides 272 Chapter 19 Nuclear Reactions 280 19.1 Introduction 280 19.2 Natural Radioactivity 280 19.3 Half-Life 282 19.4 Radioactive Series 283 19.5 Nuclear Fission and Fusion 284 19.6 Nuclear Energy 285 APPENDIX Scientific Calculations 292 A.1 Scientific Algebra 292 A.2 Calculator Mathematics 297 Glossary 312 Practice Quizzes 326 Answers to Quizzes 330 Index 335 Conversions 348 Table of the Elements 349 Periodic Table 350CHAPTER 1 Basic Concepts 1.1. INTRODUCTION Chemistry is the study of matter and energy and the interactions between them. In this chapter, we learn about the elements, which are the building blocks of every type of matter in the universe, the measurement of matter (and energy) as mass, the properties by which the types of matter can be identified, and a basic classification of matter. The symbols used to represent the elements are also presented, and an arrangement of the elements into classes having similar properties, called a periodic table, is introduced. The periodic table is invaluable to the chemist for many types of classification and understanding. Scientists have gathered so much data that they must have some way of organizing information in a useful form. Toward that end, scientific laws, hypotheses, and theories are used. These forms of generalization are introduced in Sec. 1.7. 1.2. THE ELEMENTS An element is a substance that cannot be broken down into simpler substances by ordinary means. A few more than 100 elements and the many combinations of these elements—compounds or mixtures—account for all the materials of the world. Exploration of the moon has provided direct evidence that the earth’s satellite is composed of the same elements as those on earth. Indirect evidence, in the form of light received from the sun and stars, confirms the fact that the same elements make up the entire universe. Before it was discovered on the earth, helium (from the Greek helios, meaning “sun”) was discovered in the sun by the characteristic light it emits. It is apparent from the wide variety of different materials in the world that there are a great many ways to combine elements. Changing one combination of elements to another is the chief interest of the chemist. It has long been of interest to know the composition of the crust of the earth, the oceans, and the atmosphere, since these are the only sources of raw materials for all the products that humans require. More recently, however, attention has focused on the problem of what to do with the products humans have used and no longer desire. Although elements can change combinations, they cannot be created or destroyed (except in nuclear reactions). The iron in a piece of scrap steel might rust and be changed in form and appearance, but the quantity of iron has not changed. Since there is a limited supply of available iron and since there is a limited capacity to dump unwanted wastes, recycling such materials is extremely important. The elements occur in widely varying quantities on the earth. The 10 most abundant elements make up 98% of the mass of the crust of the earth. Many elements occur only in traces, and a few elements are synthetic. Fortunately for humanity, the elements are not distributed uniformly throughout the earth. The distinct properties of the different elements cause them to be concentrated more or less, making them more available as raw materials. For example, sodium and chlorine form salt, which is concentrated in beds by being dissolved in bodies of water that later dry up. Other natural processes are responsible for the distribution of the elements that now exists on 1 Copyright © 2005, 1999, 1991 by The McGraw-Hill Companies, Inc. Click here for terms of use.2 BASIC CONCEPTS CHAP. 1 the earth. It is interesting to note that different conditions on the moon—for example, the lack of water and air on the surface—might well cause a different sort of distribution of elements on the earth’s satellite. 1.3. MATTER AND ENERGY Chemistry focusses on the study of matter, including its composition, its properties, its structure, the changes that it undergoes, and the laws governing those changes. Matter is anything that has mass and occupies space. Any material object, no matter how large or small, is composed of matter. In contrast, light, heat, and sound are forms of energy. Energy is the ability to produce change. Whenever a change of any kind occurs, energy is involved; and whenever any form of energy is changed to another form, it is evidence that a change of some kind is occurring or has occurred. The concept of mass is central to the discussion of matter and energy. The mass of an object depends on the quantity of matter in the object. The more mass the object has, the more it weighs, the harder it is to set into motion, and the harder it is to change the object’s velocity once it is in motion. Matter and energy are now known to be somewhat interconvertible. The quantity of energy producible from a quantity of matter, or vice versa, is given by Einstein’s famous equation 2 E = mc 2 where E is the energy, m is the mass of the matter that is converted to energy, and c is a constant—the square 2 of the velocity of light. The constant c is so large, 2 2 2 2 90 000 000 000 000 000 meters /second or 34 600 000 000 miles /second that tremendous quantities of energy are associated with conversions of minute quantities of matter to energy. The quantity of mass accounted for by the energy contained in a material object is so small that it is not measurable. Hence, the mass of an object is very nearly identical to the quantity of matter in the object. Particles of energy have very small masses despite having no matter whatsoever; that is, all the mass of a particle of light is associated with its energy. Even for the most energetic of light particles, the mass is small. The quantity of mass in any material body corresponding to the energy of the body is so small that it was not even conceived of until Einstein published his theory of relativity in 1905. Thereafter, it had only theoretical significance until World War II, when it was discovered how radioactive processes could be used to transform very small quantities of matter into very large quantities of energy, from which resulted the atomic and hydrogen bombs. Peaceful uses of atomic energy have developed since that time, including the production of the greater part of the electric power in many countries. The mass of an object is directly associated with its weight. The weight of a body is the pull on the body by the nearest celestial body. On earth, the weight of a body is the pull of the earth on the body, but on the moon, the weight corresponds to the pull of the moon on the body. The weight of a body is directly proportional to its mass and also depends on the distance of the body from the center of the earth or moon or whatever celestial body the object is near. In contrast, the mass of an object is independent of its position. At any given location, for example, on the surface of the earth, the weight of an object is directly proportional to its mass. When astronauts walk on the moon, they must take care to adjust to the lower gravity on the moon. Their masses are the same no matter where they are, but their weights are about one-sixth as much on the moon as on the earth because the moon is so much lighter than the earth. A given push, which would cause an astronaut to jump 1 ft high on the earth, would cause her or him to jump 6 ft on the moon. Since weight and mass are directly proportional on the surface of the earth, chemists have often used the terms interchangeably. The custom formerly was to use the term weight, but modern practice tends to use the term mass to describe quantities of matter. In this text, the term mass is used, but other chemistry texts might use the term weight, and the student must be aware that some instructors still prefer the latter. The study of chemistry is concerned with the changes that matter undergoes, and therefore chemistry is also concerned with energy. Energy occurs in many forms—heat, light, sound, chemical energy, mechanical energy, electrical energy, and nuclear energy. In general, it is possible to convert each of these forms of energy to others.CHAP. 1 BASIC CONCEPTS 3 Except for reactions in which the quantity of matter is changed, as in nuclear reactions, the law of conservation of energy is obeyed: Energy can be neither created nor destroyed (in the absence of nuclear reactions). In fact, many chemical reactions are carried out for the sole purpose of converting energy to a desired form. For example, in the burning of fuels in homes, chemical energy is converted to heat; in the burning of fuels in automobiles, chemical energy is converted to energy of motion. Reactions occurring in batteries produce electrical energy from the chemical energy stored in the chemicals from which the batteries are constructed. 1.4. PROPERTIES Every substance (Sec. 1.5) has certain characteristics that distinguish it from other substances and that may be used to establish that two specimens are the same substance or different substances. Those characteristics that serve to distinguish and identify a specimen of matter are called the properties of the substance. For example, water may be distinguished easily from iron or gold, and—although this may appear to be more difficult—iron may readily be distinguished from gold by means of the different properties of the metals. EXAMPLE 1.1. Suggest three ways in which a piece of iron can be distinguished from a piece of gold. Ans. Among other differences, 1. Iron, but not gold, will be attracted by a magnet. 2. If a piece of iron is left in humid air, it will rust. Under the same conditions, gold will undergo no appreciable change. 3. If a piece of iron and a piece of gold have exactly the same volume, the iron will have a lower mass than the gold. Physical Properties The properties related to the state (gas, liquid, or solid) or appearance of a sample are called physical properties. Some commonly known physical properties are density (Sec. 2.6), state at room temperature, color, hardness, melting point, and boiling point. The physical properties of a sample can usually be determined without changing its composition. Many physical properties can be measured and described in numerical terms, and comparison of such properties is often the best way to distinguish one substance from another. Chemical Properties A chemical reaction is a change in which at least one substance (Sec. 1.5) changes its composition and its set of properties. The characteristic ways in which a substance undergoes chemical reaction or fails to undergo chemical reaction are called its chemical properties. Examples of chemical properties are flammability, rust resistance, reactivity, and biodegradability. Many other examples of chemical properties will be presented in this book. Of the properties of iron listed in Example 1.1, only rusting is a chemical property. Rusting involves a change in composition (from iron to an iron oxide). The other properties listed do not involve any change in composition of the iron; they are physical properties. 1.5. CLASSIFICATION OF MATTER To study the vast variety of materials that exist in the universe, the study must be made in a systematic manner. Therefore, matter is classified according to several different schemes. Matter may be classified as organic or inorganic. It is organic if it is a compound of carbon and hydrogen. (A more rigorous definition of organic must wait until Chap. 18.) Otherwise, it is inorganic. Another such scheme uses the composition of matter as a basis for classification; other schemes are based on chemical properties of the various classes. For example, substances may be classified as acids, bases, or salts (Chap. 8). Each scheme is useful, allowing the study of a vast variety of materials in terms of a given class.4 BASIC CONCEPTS CHAP. 1 In the method of classification of matter based on composition, a given specimen of material is regarded as either a pure substance or a mixture. An outline of this classification scheme is shown in Table 1-1. The term pure substance (or merely substance) refers to a material all parts of which have the same composition and that has a definite and unique set of properties. In contrast, a mixture consists of two or more substances and has a somewhat arbitrary composition. The properties of a mixture are not unique, but depend on its composition. The properties of a mixture tend to reflect the properties of the substances of which it is composed; that is, if the composition is changed a little, the properties will change a little. Table 1-1 Classification of Matter Based on Composition Substances Elements Compounds Mixtures Homogeneous mixtures (solutions) Heterogeneous mixtures (mixtures) Substances There are two kinds of substances—elements and compounds. Elements are substances that cannot be broken down into simpler substances by ordinary chemical means. Elements cannot be made by the combination of simpler substances. There are slightly more than 100 elements, and every material object in the universe consists of one or more of these elements. Familiar substances that are elements include carbon, aluminum, iron, copper, gold, oxygen, and hydrogen. Compounds are substances consisting of two or more elements chemically combined in definite proportions by mass to give a material having a definite set of properties different from that of any of its constituent elements. For example, the compound water consists of 88.8% oxygen and 11.2% hydrogen by mass. The physical and chemical properties of water are distinctly different from those of both hydrogen and oxygen. For example, water is a liquid at room temperature and pressure, while the elements of which it is composed are gases under these same conditions. Chemically, water does not burn; hydrogen may burn explosively in oxygen (or air). Any sample of pure water, regardless of its source, has the same composition and the same properties. There are millions of known compounds, and thousands of new ones are discovered or synthesized each year. Despite such a vast number of compounds, it is possible for the chemist to know certain properties of each one, because compounds can be classified according to their composition and structure, and groups of compounds in each class have some properties in common. For example, organic compounds are generally combustible in excess oxygen, yielding carbon dioxide and water. So any compound that contains carbon and hydrogen may be predicted by the chemist to be combustible in oxygen. Organic compound + oxygen −→ carbon dioxide + water + possible other products Mixtures There are two kinds of mixtures—homogeneous mixtures and heterogeneous mixures. Homogeneous mix- tures are also called solutions, and heterogeneous mixtures are sometimes simply called mixtures. In heteroge- neous mixtures, it is possible to see differences in the sample merely by looking, although a microscope might be required. In contrast, homogeneous mixtures look the same throughout the sample, even under the best optical microscope. EXAMPLE 1.2. A teaspoon of salt is added to a cup of warm water. White crystals are seen at the bottom of the cup. Is the mixture homogeneous or heterogeneous? Then the mixture is stirred until the salt crystals disappear. Is the mixture now homogeneous or heterogeneous? Ans. Before stirring, the mixture is heterogeneous; after stirring, the mixture is homogeneous—a solution.CHAP. 1 BASIC CONCEPTS 5 Distinguishing a Mixture from a Compound Let us imagine an experiment to distinguish a mixture from a compound. Powdered sulfur is yellow and it dissolves in carbon disulfide, but it is not attracted by a magnet. Iron filings are black and are attracted by a magnet, but do not dissolve in carbon disulfide. You can mix iron filings and powdered sulfur in any ratio and get a yellowish-black mixture—the more sulfur that is present, the yellower the mixture will be. If you put the mixture in a test tube and hold a magnet alongside the test tube just above the mixture, the iron filings will be attracted, but the sulfur will not. If you pour enough (colorless) carbon disulfide on the mixture, stir, and then pour off the resulting yellow liquid, the sulfur dissolves but the iron does not. The mixture of iron filings and powdered sulfur is described as a mixture because the properties of the combination are still the properties of its components. If you mix sulfur and iron filings in a certain proportion and then heat the mixture, you can see a red glow spread through the mixture. After it cools, the black solid lump that is produced—even if crushed into a powder—does not dissolve in carbon disulfide and is not attracted by a magnet. The material has a new set of properties; it is a compound, called iron(II) sulfide. It has a definite composition; and if, for example, you had mixed more iron with the sulfur originally, some iron(II) sulfide and some leftover iron would have resulted. The extra iron would not have become part of the compound. 1.6. REPRESENTATION OF ELEMENTS Each element has an internationally accepted symbol to represent it. A list of the names and symbols of the elements is found on page 349 of this book. Note that symbols for the elements are for the most part merely abbreviations of their names, consisting of either one or two letters. The first letter of the symbol is always written as a capital letter; the second letter, if any, is always written as a lowercase (small) letter. The symbols of a few elements do not suggest their English names, but are derived from the Latin or German names of the elements. The 10 elements whose names do not begin with the same letter as their symbols are listed in Table 1-2. For convenience, on page 349 of this book, these elements are listed twice—once alphabetically by name and again under the letter that is the first letter of their symbol. It is important to memorize the names and symbols of the most common elements. To facilitate this task, the most familiar elements are listed in Table 1-3. The elements with symbols in bold type should be learned first. Table 1-2 Symbols and Names with Different Initials Symbol Name Symbol Name Ag Silver Na Sodium Au Gold Pb Lead Fe Iron Sb Antimony Hg Mercury Sn Tin K Potassium W Tungsten The Periodic Table A convenient way of displaying the elements is in the form of a periodic table, such as is shown on page 350 of this book. The basis for the arrangement of elements in the periodic table will be discussed at length in Chaps. 3 and 4. For the present, the periodic table is regarded as a convenient source of general information about the elements. It will be used repeatedly throughout the book. One example of its use, shown in Fig. 1-1, is to classify the elements as metals or nonmetals. All the elements except hydrogen that lie to the left of the stepped line drawn on the periodic table, starting to the left of B and descending stepwise to a point between Po and At, are metals. The other elements are nonmetals. It is readily seen that the majority of elements are metals.6 BASIC CONCEPTS CHAP. 1 Table 1-3 Important Elements Whose Names and Symbols Should Be Known 1 2 H He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 46 47 48 50 51 52 53 54 Rb Sr Pd Ag Cd Sn Sb Te I Xe 55 56 74 78 79 80 82 83 86 Cs Ba W Pt Au Hg Pb Bi Rn 92 U The smallest particle of an element that retains the composition of the element is called an atom. Details of the nature of atoms are given in Chaps. 3 and 4. The symbol of an element is used to stand for one atom of the element as well as for the element itself. 1.7. LAWS, HYPOTHESES, AND THEORIES In chemistry, as in all sciences, it is necessary to express ideas in terms having very precise meanings. These meanings are often unlike the meanings of the same words in nonscientific usage. For example, the meaning of the word property as used in chemistry can be quite different from its meaning in ordinary conversation. Also, in chemical terminology, a concept may be represented by abbreviations, such as symbols or formulas, or by some mathematical expression. Together with precisely defined terms, such symbols and mathematical expressions constitute a language of chemistry. This language must be learned well. As an aid to recognition of special terms, when such terms are used for the first time in this book, they will be italicized. A statement that generalizes a quantity of experimentally observable phenomena is called a scientific law. For example, if a person drops a pencil, it falls downward. This result is predicted by the law of gravity. A generalization that attempts to explain why certain experimental results occur is called a hypothesis. When a hypothesis is accepted as true by the scientific community, it is then called a theory. One of the most important scientific laws is the law of conservation of mass: During any process (chemical reaction, physical change, or Nonmetal B Al Si Nonmetals Ge As Sb Te Po At Metals Fig. 1-1. Metals and nonmetalsCHAP. 1 BASIC CONCEPTS 7 even a nuclear reaction) mass is neither created nor destroyed. Because of the close approximation that the mass of an object is the quantity of matter it contains (excluding the mass corresponding to its energy) the law of conservation of mass can be approximated by the law of conservation of matter: During an ordinary chemical reaction, matter can be neither created nor destroyed. EXAMPLE 1.3. When a piece of iron is left in moist air, its surface gradually turns brown and the object gains mass. Explain this phenomenon. Ans. The brown material is an iron oxide, rust, formed by a reaction of the iron with the oxygen in the air. Iron + oxygen −→ an iron oxide The increase in mass is just the mass of the combined oxygen. When a long burns, the ash (which remains) is much lighter than the original log, but this is not a contradiction of the law of conservation of matter. In addition to the log, which consists mostly of compounds containing carbon, hydrogen, and oxygen, oxygen from the air is consumed by the reaction. In addition to the ash, carbon dioxide and water vapor are produced by the reaction. Log + oxygen −→ ash + carbon dioxide + water vapor The total mass of the ash plus the carbon dioxide and the water vapor is equal to the total mass of the log plus the oxygen. As always, the law of conservation of matter is obeyed as precisely as chemists can measure. The law of conservation of mass is fundamental to the understanding of chemical reactions. Other laws related to the behavior of matter are equally important, and learning how to apply these laws correctly is a necessary goal of the study of chemistry. Solved Problems 1.1. Are elements heterogeneous or homogeneous? Ans. Homogeneous. They look alike throughout the sample because they are alike throughout the sample. 1.2. How can you tell if the word mixture means any mixture or a heterogeneous mixture? Ans. You can tell from the context. For example, if a problem asks if a sample is a solution or a mixture, the word mixture means heterogeneous mixture. If it asks whether the sample is a compound or a mixture, it means any kind of mixture. Such usage occurs in ordinary English as well as in technical usage. For example, the word day has two meanings—one is a subdivision of the other. ”How many hours are there in a day? What is the opposite of night?” 1.3. Are compounds heterogeneous or homogeneous? Ans. Homogeneous. They look alike throughout the sample because they are alike throughout the sample. Since there is only one substance present, even if it is a combination of elements, it must be alike throughout. 1.4. A generality states that all compounds containing both carbon and hydrogen burn. Do octane and propane burn? (Each contains only carbon and hydrogen.) Ans. Yes, both burn. It is easier to learn that all organic compounds burn than to learn a list of millions of organic compounds that burn. On an examination, however, a question will probably specify one particular organic compound. You must learn a generality and be able to respond to a specific example of it. 1.5. Sodium is a very reactive metallic element; for example, it liberates hydrogen gas when treated with water. Chlorine is a yellow-green, choking gas, used in World War I as a poison gas. Contrast these properties with those of the compound of sodium and chlorine—sodium chloride—known as table salt. Ans. Salt does not react with water to liberate hydrogen, is not reactive, and is not poisonous. It is a white solid and not a silvery metal or a green gas. In short, it has its own set of properties; it is a compound.8 BASIC CONCEPTS CHAP. 1 1.6. TNT is a compound of carbon, nitrogen, hydrogen, and oxygen. Carbon occurs is two common forms— graphite (the material in “lead pencils”) and diamond. Oxygen and nitrogen comprise over 98% of the atmosphere. Hydrogen is an element that reacts explosively with oxygen. Which of the properties of the elements determines the properties of TNT? Ans. The properties of the elements do not matter. The properties of the compound are quite independent of those of the elements. A compound has its own distinctive set of properties. TNT is most noted for its explosiveness. 1.7. What properties of stainless steel make it more desirable for many purposes than ordinary steel? Ans. Its resistance to rusting and corrosion. 1.8. What properties of DDT make it useful? What properties make it undesirable? Ans. DDT’s toxicity to insects is its useful property; its toxicity to humans, birds, and other animals makes it undesirable. It is stable, that is, nonbiodegradable (does not decompose spontaneously to simpler substances in the environment). This property makes its use as an insecticide more difficult. 1.9. Name an object or an instrument that changes: (a) chemical energy to heat (d ) electrical energy to light (b) chemical energy to electrical energy (e) motion to electrical energy (c) electrical energy to chemical energy (f ) electrical energy to motion Ans. (a) gas stove (d ) lightbulb (b) battery (e) generator or alternator (c) rechargeable battery (f ) electric motor 1.10. A sample contains 88.8% oxygen and 11.2% hydrogen by mass, and is gaseous and explosive at room temperature. (a) Is the sample a compound or a mixture? (b) After the sample explodes and cools, it is a liquid. Is the sample now a compound or a mixture? (c) Would it be easier to change the percentage of oxygen before or after the explosion? Ans. (a) The sample is a mixture. (The compound of hydrogen and oxygen with this composition—water—is a liquid under these conditions.) (b) It is a compound. (c) Before the explosion. It is easy to add hydrogen or oxygen to the gaseous mixture, but you cannot change the composition of water. 1.11. Name one exception to the statement that nonmetals lie to the right of the stepped line in the periodic table (page 350). Ans. Hydrogen 1.12. Calculate the ratio of the number of metals to the number of nonmetals in the periodic table (page 350). Ans. There are 109 elements whose symbols are presented, of which 22 are nonmetals and 87 are metals, so the ratio is 3.95 metals per nonmetal. 1.13. Give the symbol for each of the following elements: (a) iron, (b) copper, (c) carbon, (d ) sodium, (e) silver, ( f ) aluminum. Ans. (a)Fe (b)Cu (c)C (d)Na (e)Ag (f)Al 1.14. Name each of the following elements: (a)K,(b)P,(c) Cl, (d)H,(e)O. Ans. (a) Potassium (b) Phosphorus (c) Chlorine (d ) Hydrogen (e) OxygenCHAP. 1 BASIC CONCEPTS 9 1.15. Distinguish between a theory and a law. Ans. A law tells what happens under a given set of circumstances, while a theory attempts to explain why that behavior occurs. 1.16. Distinguish clearly between (a) mass and matter and (b) mass and weight. Ans. (a) Matter is any kind of material. The mass of an object depends mainly on the matter that it contains. It is affected only very slightly by the energy in it. (b) Weight is the attraction of the earth on an object. It depends on the mass of the object and its distance to the center of the earth.

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