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List of the Elements with Their Atomic Symbols and Atomic Weights Name Symbol Atomic Atomic Number Weight Actinium Ac 89 (227)* Aluminum Al 13 26.981538 Americium Am 95 (243) Antimony Sb 51 121.760 Argon Ar 18 39.948 Arsenic As 33 74.92160 Astatine At 85 (210) Barium Ba 56 137.327 Berkelium Bk 97 (247) Beryllium Be 4 9.012182 Bismuth Bi 83 208.98040 Bohrium Bh 107 (272) Boron B 5 10.811 Bromine Br 35 79.904 Cadmium Cd 48 112.411 Calcium Ca 20 40.078 Californium Cf 98 (251) Carbon C 12.0107 Cerium Ce 58 140.116 Cesium Cs 55 132.90545 Chlorine Cl 17 35.453 Chromium Cr 24 51.9961 Cobalt Co 27 58.933195 Copernicium Cn 112 (285) Copper Cu 29 63.546 Curium Cm 96 (247 ) Darmstadtium Ds 110 (281) Dubnium Db 105 (268) Dysprosium Dy 66 162.500 Einsteinium Es 99 (252) Erbium Er 68 167.259 Europium Eu 63 151.964 Fermium Fm 100 (257) Flerovium Fl 114 (289) Fluorine F 9 18.998403 Francium Fr 87 (223) Gadolinium Gd 64 157.25 Gallium Ga 31 69.723 Germanium Ge 32 72.64 Gold Au 79 196.96657 Hafnium Hf 72 178.49 Hassium Hs 108 (270) a Helium He 2 4.002602 Holmium Ho 67 164.93032 Hydrogen H 1 1.00794 Indium In 49 114.818 Iodine I 53 126.90447 Iridium Ir 77 192.217 Iron Fe 26 55.845 Krypton Kr 36 83.798 Lanthanum La 57 138.9055 Lawrencium Lr 103 (262) Lead Pb 82 207.2 Lithium Li 3 6.941 Livermorium Lv 116 (293) Lutetium Lu 71 174.9668 Magnesium Mg 12 24.3050 Manganese Mn 25 54.938045 Meitnerium Mt 109 (276) Atomic Atomic Name Symbol Number Weight Mendelevium Md 101 (258) Mercury Hg 80 200.59 Molybdenum Mo 42 95.96 Moscovium Mc 115 (288) Neodymium Nd 60 144.242 Neon Ne 10 20.1797 Neptunium Np 93 (237) Nickel Ni 28 58.6934 Nihonium Nh 113 (284) Niobium Nb 41 92.90638 Nitrogen N 7 14.0067 Nobelium No 102 (259) Oganesson Og 118 (294) Osmium Os 76 190.23 Oxygen O 8 15.9994 Palladium Pd 46 106.42 Phosphorus P 15 30.973762 Platinum Pt 78 195.094 Plutonium Pu 94 (244) Polonium Po 84 (209) Potassium K 19 39.0983 Praseodymium Pr 59 140.90765 Promethium Pm 61 (145) Protactinium Pa 91 231.03588 Radium Ra 88 (226) Radon Rn 86 (222) a Rhenium Re 75 186.207 Rhodium Rh 45 102.90550 Roentgenium Rg 111 (280) Rubidium Rb 37 85.4678 Ruthenium Ru 44 101.07 Rutherfordium Rf 104 (265) Samarium Sm 62 150.36 Scandium Sc 21 44.955912 Seaborgium Sg 106 (271) Selenium Se 34 78.96 Silicon Si 14 28.0855 Silver Ag 47 107.8682 Sodium Na 11 22.989769 Strontium Sr 38 87.62 Sulfur S 16 32.065 Tantalum Ta 73 180.9479 Technetium Tc 43 (98) Tellurium Te 52 127.60 Tennessine Ts 117 (292) Terbium Tb 65 158.92535 Thallium Tl 81 204.3833 Thorium Th 90 232.0381 Thulium Tm 69 168.93421 Tin Sn 50 118.710 Titanium Ti 22 47.867 Tungsten W 74 183.84 Uranium U 92 238.02891 Vanadium V 23 50.9415 Xenon Xe 54 131.293 Ytterbium Yb 70 173.054 Yttrium Y 39 88.90585 Zinc Zn 30 65.38 Zirconium Zr 40 91.224 *Values in parentheses are the mass numbers of the most common or longest lived isotopes of radioactive elements CVR_MCMU6230_08_SE_FEP.indd 04/12/2018 09:47 CVR_MCMU6230_08_SE_FEP.indd 04/12/2018 09:47 137.327 88 Ra (226) 87 Fr (223) (265) 57 La (262) Lanthanide series Actinide series 58 Ce 104 Rf 103 Lr (227) 89 Ac 138.9055 (268) 178.49 174.9668 59 Pr (271) 106 Sg 183.84 91 Pa 92 U 144.242 60 Nd (272) 107 Bh 186.207 75 Re (98) 232.0381 231.03588 238.02891 90 Th 140.116 140.90765 105 Db 180.9479 74 W 95.96 (237) 93 Np (145) 61 Pm (270) 108 Hs 190.23 76 Os 101.07 44 Ru (244) 94 Pu 150.36 62 Sm (276) 109 Mt 192.217 77 Ir 102.90550 45 Rh (243) 95 Am 151.964 63 Eu (281) 110 Ds 195.094 78 Pt 106.42 46 Pd 58.933195 58.6934 (247) 96 Cm 157.25 64 Gd (280) 111 Rg 196.96657 79 Au 107.8682 47 Ag 63.546 66 Dy (284) 113 Nh 204.3833 81 Tl 114.818 49 In 69.723 31 Ga 67 Ho (289) 114 FL 207.2 82 Pb 118.710 50 Sn 72.64 32 Ge 68 Er (288) 115 Mc 208.98040 83 Bi 121.760 51 Sb 74.92160 33 As 26.981538 28.0855 30.973762 15 P 14.0067 N 15 5A F 17 7A (247) 97 Bk (251) 98 Cf (252) 99 Es (257) 100 Fm 10 Ne 4.002602 He 18 8A 69 Tm (293) 116 Lv (209) 84 Po 127.60 52 Te 78.96 34 Se 32.065 16 S (258) 101 Md 54 Xe 83.798 36 Kr 39.948 18 Ar 70 Yb (292) 117 Ts (210) 85 At (259) 102 No (294) 118 Og (222) 86 Rn 126.90447 131.293 53 I 79.904 35 Br 35.453 17 Cl 15.9994 18.998403 20.1797 O 16 6A Main groups 158.92535 162.500 164.93032 167.259 168.93421 173.054 65 Tb (285) 112 Cn 200.59 80 Hg 112.411 48 Cd 65.38 30 Zn 132.90545 73 Ta 72 Hf 71 Lu 92.90638 91.224 88.90585 43 Tc 29 Cu 56 Ba 42 Mo 28 Ni 87.62 26 Fe 51.9961 54.938045 55.845 25 Mn 27 Co 55 Cs 40 Zr 39 Y 24 Cr 85.4678 50.9415 47.867 44.955912 41 Nb 23 V 22 Ti 21 Sc 38 Sr 12 2B 40.078 11 1B 37 Rb 10 39.0983 8B 20 Ca 24.3050 7B 19 K 6B 22.989769 5B 4B 3B 14 Si 11 Na 12.0107 12 Mg 6.941 13 Al 9.012182 Li 10.811 C B Be 1.00794 Transition metals 14 4A 13 3A 2A H 1A Main groups Periodic Table of the Elements CHEMISTRY E I G H T H JILL K ROBINSON Indiana University JOHN E MCMURRY Cornell University ROBERT C FAY Cornell University E D I T I O N Director of Portfolio Management: Jeanne Zalesky Executive Courseware Portfolio Manager: Terry Haugen Content Producer: Shercian Kinosian Managing Producer: Kristen Flathman Courseware Director, Content Development: Barbara Yien Courseware Analysts: Cathy Murphy, Coleen Morrison, Jay McElroy Courseware Editorial Assistant: Harry Misthos Rich Media Content Producers: Jenny Moryan, Ziki Dekel Director MasteringChemistry Content Development: Amir Said MasteringChemistry Senior Content Producer: Margaret Trombley MasteringChemistry Content Producers: Meaghan Fallano, Kaitlin Smith Full-Service Vendor, Project Manager: Pearson CSC, Kelly Murphy Copyeditor: Pearson CSC Compositor: Pearson CSC Art House, Coordinator: Lachina, Rebecca Marshall Design Manager: Maria Guglielmo Walsh Interior & Cover Designer: Gary Hespeneide Rights & Permissions Manager: Ben Ferrini Rights & Permissions Project Manager: Pearson CSC, Eric Schrader Rights & Permissions Specialist/Photo Researcher: Pearson CSC, Angelica Aranas Manufacturing Buyer: Stacey Weinberger VP, Director of Field Marketing: Tim Galligan Director of Product Marketing: Allison Rona Executive Field Marketing Manager: Christopher Barker Senior Product Marketing Manager: Elizabeth Bell Cover Photo Credit: Beauty of Science/Science Source Copyright © 2020, 2016, 2012 by Pearson Education, Inc 221 River Street, Hoboken, NJ 07030 All Rights Reserved Printed in the United States of America This publication is protected by copyright, and permission should be obtained from the publisher prior to any prohibited reproduction, storage in a retrieval system, or transmission in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise For information regarding permissions, request forms and the appropriate contacts within the Pearson Education Global Rights & Permissions department Attributions of third party content appear on page C-1, which constitutes an extension of this copyright page PEARSON, ALWAYS LEARNING, Mastering™ Chemistry, and Learning Catalytics™ are exclusive trademarks in the U.S and/or other countries owned by Pearson Education, Inc or its affiliates Unless otherwise indicated herein, any third-party trademarks that may appear in this work are the property of their respective owners and any references to third-party trademarks, logos or other trade dress are for demonstrative or descriptive purposes only Such references are not intended to imply any sponsorship, endorsement, authorization, or promotion of Pearson's products by the owners of such marks, or any relationship between the owner and Pearson Education, Inc or its affiliates, authors, licensees or distributors Library of Congress Cataloging-in-Publication Data Names: Robinson, Jill K | McMurry, John | Fay, Robert C., 1936Title: Chemistry / Jill K Robinson (Indiana University), John E McMurry (Cornell University), Robert C Fay (Cornell University) Description: Eighth edition | Hoboken, NJ : Pearson Education, Inc., [2020] Identifiers: LCCN 2018053050 | ISBN 9780134856230 (casebound) Subjects: LCSH: Chemistry Textbooks Classification: LCC QD33.2 M36 2020 | DDC 540 dc23 LC record available at https://lccn.loc.gov/2018053050 1 19 www.pearson.com ISBN 10: 0-134-85623-6 ISBN 13: 978-0-134-85623-0 (Student edition) ISBN 10: 0-135-21012-7 ISBN 13:978-0-135-21012-3 (Looseleaf Edition) Brief Contents Preface xiii For Instructors xvi 1 Chemical Tools: Experimentation and Measurement 2 Atoms, Molecules, and Ions 33 3 Mass Relationships in Chemical Reactions 83 4 Reactions in Aqueous Solution 116 5 Periodicity and the Electronic Structure of Atoms 161 6 Ionic Compounds: Periodic Trends and Bonding Theory 208 7 Covalent Bonding and Electron-Dot Structures 238 8 Covalent Compounds: Bonding Theories and Molecular Structure 278 9 Thermochemistry: Chemical Energy 327 10 Gases: Their Properties and Behavior 374 11 Liquids and Phase Changes 422 12 Solids and Solid-State Materials 450 13 Solutions and Their Properties 494 14 Chemical Kinetics 538 15 Chemical Equilibrium 601 16 Aqueous Equilibria: Acids and Bases 654 17 Applications of Aqueous Equilibria 708 18 Thermodynamics: Entropy, Free Energy, and Spontaneity 768 19 Electrochemistry 813 20 Nuclear Chemistry 870 21 Transition Elements and Coordination Chemistry 904 22 The Main-Group Elements 954 23 Organic and Biological Chemistry 1003 iii Contents Preface xiii For Instructors xvi 2.12 Ions and Ionic Bonds 61 2.13 Naming Chemical Compounds 63 INQUIRY Chemical Tools: Experimentation and Measurement 1 The Scientific Method: Nanoparticle Catalysts for Fuel Cells 2 1.2 Measurements: SI Units and Scientific Notation 1.3 Mass and Its Measurement 1.4 Length and Its Measurement 1.5 Temperature and Its Measurement 1.6 Derived Units: Volume and Its Measurement 11 1.7 Derived Units: Density and Its Measurement 13 1.8 Derived Units: Energy and Its Measurement 14 1.9 Accuracy, Precision, and Significant Figures in Measurement 16 1.10 Significant Figures in Calculations 18 1.11 Converting from One Unit to Another 20 1.1 INQUIRY hat are the unique properties of nanoscale W materials? 23 Study Guide • Key Terms • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Mass Relationships in Chemical Reactions 83 3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 3.9 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Chemistry and the Elements 34 Elements and the Periodic Table 36 Some Common Groups of Elements and Their Properties 38 2.4 Observations Supporting Atomic Theory: The Conservation of Mass and the Law of Definite Proportions 41 2.5 The Law of Multiple Proportions and Dalton’s Atomic Theory 43 2.6 Atomic Structure: Electrons 45 2.7 Atomic Structure: Protons and Neutrons 47 2.8 Atomic Numbers 49 2.9 Atomic Weights and the Mole 51 2.10 Measuring Atomic Weight: Mass Spectrometry 55 2.11 Mixtures and Chemical Compounds; Molecules and Covalent Bonds 57 Representing Chemistry on Different Levels 84 Balancing Chemical Equations 85 Molecular Weight and Molar Mass 88 Stoichiometry: Relating Amounts of Reactants and Products 90 Yields of Chemical Reactions 92 Reactions with Limiting Amounts of Reactants 94 Percent Composition and Empirical Formulas 97 Determining Empirical Formulas: Elemental Analysis 100 Determining Molecular Weights: Mass Spectrometry 103 INQUIRY ow is the principle of atom economy H used to minimize waste in a chemical synthesis? 105 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Atoms, Molecules, and Ions 33 2.1 2.2 2.3 ow can measurements of oxygen H and hydrogen isotopes in ice cores determine past climates? 69 Reactions in Aqueous Solution 116 4.1 4.2 4.3 4.4 4.5 4.6 4.7 4.8 4.9 Solution Concentration: Molarity 117 Diluting Concentrated Solutions 119 Electrolytes in Aqueous Solution 121 Types of Chemical Reactions in Aqueous Solution 123 Aqueous Reactions and Net Ionic Equations 124 Precipitation Reactions and Solubility Guidelines 125 Acids, Bases, and Neutralization Reactions 128 Solution Stoichiometry 132 Measuring the Concentration of a Solution: Titration 133 4.10 4.11 4.12 4.13 4.14 Contents Oxidation–Reduction (Redox) Reactions 135 Identifying Redox Reactions 138 The Activity Series of the Elements 141 Redox Titrations 144 Some Applications of Redox Reactions 146 INQUIRY Periodicity and the Electronic Structure of Atoms 161 5.2 5.3 5.4 5.5 5.6 5.7 5.8 5.9 5.10 5.11 5.12 5.13 Wave Properties of Radiant Energy and the Electromagnetic Spectrum 162 Particlelike Properties of Radiant Energy: The Photoelectric Effect and Planck’s Postulate 166 Atomic Line Spectra and Quantized Energy 169 Wavelike Properties of Matter: de Broglie’s Hypothesis 173 The Quantum Mechanical Model of the Atom: Heisenberg’s Uncertainty Principle 175 The Quantum Mechanical Model of the Atom: Orbitals and Quantum Numbers 176 The Shapes of Orbitals 179 Electron Spin and the Pauli Exclusion Principle 184 Orbital Energy Levels in Multielectron Atoms 185 Electron Configurations of Multielectron Atoms 187 Anomalous Electron Configurations 189 Electron Configurations and the Periodic Table 189 Electron Configurations and Periodic Properties: Atomic Radii 192 INQUIRY ow does knowledge of atomic emission H spectra help us build more efficient light bulbs? 195 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Ionic Compounds: Periodic Trends and Bonding Theory 208 6.1 6.2 6.3 6.4 Electron Configurations of Ions 209 Ionic Radii 212 Ionization Energy 214 Higher Ionization Energies 216 Electron Affinity 218 The Octet Rule 220 Ionic Bonds and the Formation of Ionic Solids 222 Lattice Energies in Ionic Solids 226 INQUIRY ow sports drinks replenish H the substances lost in sweat? 148 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 5.1 6.5 6.6 6.7 6.8 v ow ionic liquids lead to more H environmentally friendly processes? 228 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Covalent Bonding and Electron-Dot Structures 238 Covalent Bonding in Molecules 239 Strengths of Covalent Bonds 240 Polar Covalent Bonds: Electronegativity 242 A Comparison of Ionic and Covalent Compounds 246 Electron-Dot Structures: The Octet Rule 247 Procedure for Drawing Electron-Dot Structures 250 Drawing Electron-Dot Structures for Radicals 254 Electron-Dot Structures of Compounds Containing Only Hydrogen and Second-Row Elements 255 7.9 Electron-Dot Structures and Resonance 257 7.10 Formal Charges 261 7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8 INQUIRY ow does bond polarity affect the toxicity H of organophosphate insecticides? 265 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Covalent Compounds: Bonding Theories and Molecular Structure 278 8.1 8.2 8.3 8.4 8.5 8.6 8.7 8.8 8.9 Molecular Shapes: The VSEPR Model 279 Valence Bond Theory 286 Hybridization and sp3 Hybrid Orbitals 287 Other Kinds of Hybrid Orbitals 290 Polar Covalent Bonds and Dipole Moments 295 Intermolecular Forces 298 Molecular Orbital Theory: The Hydrogen Molecule 306 Molecular Orbital Theory: Other Diatomic Molecules 308 Combining Valence Bond Theory and Molecular Orbital Theory 312 INQUIRY hich is better for human health, natural or W synthetic vitamins? 314 Study Guide • Key Terms • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems vi Contents Thermochemistry: Chemical Energy 327 Energy and Its Conservation 328 Internal Energy and State Functions 330 Expansion Work 332 Energy and Enthalpy 334 Thermochemical Equations and the Thermodynamic Standard State 336 9.6 Enthalpies of Chemical and Physical Changes 338 9.7 Calorimetry and Heat Capacity 341 9.8 Hess’s Law 345 9.9 Standard Heats of Formation 348 9.10 Bond Dissociation Energies 350 9.11 An Introduction to Entropy 352 9.12 An Introduction to Free Energy 355 11.4 Energy Changes during Phase Transitions 431 11.5 Phase Diagrams 433 11.6 Liquid Crystals 436 9.1 9.2 9.3 9.4 9.5 INQUIRY ow we determine the energy content H of biofuels? 359 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 10 Gases: Their Properties and Behavior 374 Gases and Gas Pressure 375 The Gas Laws 380 The Ideal Gas Law 385 Stoichiometric Relationships with Gases 387 Mixtures of Gases: Partial Pressure and Dalton’s Law 390 10.6 The Kinetic–Molecular Theory of Gases 393 10.7 Gas Diffusion and Effusion: Graham’s Law 395 10.8 The Behavior of Real Gases 397 10.9 The Earth’s Atmosphere and the Greenhouse Effect 398 10.10 Greenhouse Gases 401 10.11 Climate Change 403 10.1 10.2 10.3 10.4 10.5 INQUIRY How inhaled anesthetics work? 407 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 11 Liquids and Phase Changes 422 11.1 Properties of Liquids 423 11.2 Vapor Pressure and Boiling Point 424 11.3 Phase Changes between Solids, Liquids, and Gases 428 INQUIRY How is caffeine removed from coffee? 439 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 12 Solids and Solid-State Materials 450 12.1 Types of Solids 451 12.2 Probing the Structure of Solids: X-Ray Crystallography 453 12.3 The Packing of Spheres in Crystalline Solids: Unit Cells 455 12.4 Structures of Some Ionic Solids 459 12.5 Structures of Some Covalent Network Solids 462 12.6 Bonding in Metals 464 12.7 Semiconductors 468 12.8 Semiconductor Applications 471 12.9 Superconductors 475 12.10 Ceramics and Composites 477 INQUIRY hat are quantum dots, and what controls W their color? 482 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 13 Solutions and Their Properties 494 Solutions 495 Enthalpy Changes and the Solution Process 496 Predicting Solubility 498 Concentration Units for Solutions 501 Some Factors That Affect Solubility 506 Physical Behavior of Solutions: Colligative Properties 510 13.7 Vapor-Pressure Lowering of Solutions: Raoult’s Law 511 13.8 Boiling-Point Elevation and Freezing-Point Depression of Solutions 517 13.9 Osmosis and Osmotic Pressure 521 13.1 13.2 13.3 13.4 13.5 13.6 INQUIRY ow does hemodialysis cleanse the blood H of patients with kidney failure? 525 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems Contents 14 Chemical Kinetics 538 14.1 Reaction Rates 539 14.2 Rate Laws and Reaction Order 544 14.3 Method of Initial Rates: Experimental Determination of a Rate Law 546 14.4 Integrated Rate Law: Zeroth-Order Reactions 550 14.5 Integrated Rate Law: First-Order Reactions 552 14.6 Integrated Rate Law: Second-Order Reactions 557 14.7 Reaction Rates and Temperature: The Arrhenius Equation 560 14.8 Using the Arrhenius Equation 564 14.9 Reaction Mechanisms 567 14.10 Rate Laws for Elementary Reactions 570 14.11 Rate Laws for Overall Reactions 573 14.12 Catalysis 577 14.13 Homogeneous and Heterogeneous Catalysts 580 INQUIRY How enzymes work? 583 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 15 Chemical Equilibrium 601 The Equilibrium State 603 The Equilibrium Constant Kc 605 The Equilibrium Constant KP 610 Heterogeneous Equilibria 612 Using the Equilibrium Constant 614 Factors That Alter the Composition of an Equilibrium Mixture: Le Châtelier’s Principle 624 15.7 Altering an Equilibrium Mixture: Changes in Concentration 625 15.8 Altering an Equilibrium Mixture: Changes in Pressure and Volume 629 15.9 Altering an Equilibrium Mixture: Changes in Temperature 631 15.10 The Link between Chemical Equilibrium and Chemical Kinetics 634 15.1 15.2 15.3 15.4 15.5 15.6 INQUIRY ow does high altitude affect oxygen H transport in the body? 637 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems vii Dissociation of Water 664 The pH Scale 666 Measuring pH 668 The pH in Solutions of Strong Acids and Strong Bases 669 16.8 Equilibria in Solutions of Weak Acids 671 16.9 Calculating Equilibrium Concentrations in Solutions of Weak Acids 673 16.10 Percent Dissociation in Solutions of Weak Acids 677 16.11 Polyprotic Acids 678 16.12 Equilibria in Solutions of Weak Bases 682 16.13 Relation Between Ka and Kb 684 16.14 Acid–Base Properties of Salts 686 16.15 Lewis Acids and Bases 691 16.4 16.5 16.6 16.7 INQUIRY as the problem of acid rain been H solved? 694 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 17 Applications of Aqueous Equilibria 708 17.1 Neutralization Reactions 709 17.2 The Common-Ion Effect 712 17.3 Buffer Solutions 716 17.4 The Henderson–Hasselbalch Equation 720 17.5 pH Titration Curves 723 17.6 Strong Acid–Strong Base Titrations 724 17.7 Weak Acid–Strong Base Titrations 727 17.8 Weak Base–Strong Acid Titrations 732 17.9 Polyprotic Acid–Strong Base Titrations 733 17.10 Solubility Equilibria for Ionic Compounds 738 17.11 Measuring Ksp and Calculating Solubility from Ksp 739 17.12 Factors That Affect Solubility 742 17.13 Precipitation of Ionic Compounds 750 17.14 Separation of Ions by Selective Precipitation 751 17.15 Qualitative Analysis 752 INQUIRY What is causing ocean acidification? 754 Study Guide • Key Terms • Key Equations • Practice Test • Conceptual Problems • Section Problems • Multiconcept Problems 16 Aqueous Equilibria: Acids and Bases 654 18 Thermodynamics: Entropy, Free Energy, and Spontaneity 768 16.1 Acid–Base Concepts: The Brønsted–Lowry Theory 655 16.2 Acid Strength and Base Strength 658 16.3 Factors That Affect Acid Strength 661 18.1 Spontaneous Processes 769 18.2 Enthalpy, Entropy, and Spontaneous Processes 770 18.3 Entropy and Probability 773 2.11 Mixtures and Chemical Compounds; Molecules and Covalent Bonds 59 CONCEPTUAL WORKED EXAMPLE 2.8 Visual Representations of Mixtures and Compounds Which of the following drawings represents a mixture, which a pure compound, and which an element? (a) (b) (c) STRATEGY Most people (professional chemists included) find chemistry easier to grasp when they can visualize the behavior of atoms, thereby turning symbols into pictures The Conceptual Problems in this text are intended to help you that, frequently representing atoms and molecules as collections of spheres Don’t take the pictures literally; focus instead on interpreting what they represent An element contains only one kind of atom, while a compound contains two or more different elements bonded together A pure substance contains only one type of element or compound, while a mixture contains two or more substances ▶▶CONCEPTUAL APPLY 2.16 Red and blue spheres represent atoms of different elements (a) Which drawing(s) illustrate a pure substance? (b) Which drawing(s) illustrate a mixture? (c) Which two drawings illustrate the law of multiple proportions? (a) (b) (c) (d) SOLUTION Drawing (a) represents a mixture of two diatomic elements, one composed of two red atoms and one composed of two blue atoms Drawing (b) represents molecules of a pure diatomic element because all atoms are identical Drawing (c) represents molecules of a pure compound composed of one red and one blue atom ▶▶CONCEPTUAL PRACTICE 2.15 Which of the following drawings represents a pure sample of hydrogen peroxide (H2O2) molecules? The red spheres represent oxygen atoms, and the ivory spheres represent hydrogen (a) (b) (c) (d) 60 chapter Atoms, Molecules, and Ions Chemists normally represent a molecule by giving its structural formula, which shows the specific connections between atoms and therefore gives much more information than the chemical formula alone Ethyl alcohol, for example, has the chemical formula C2H6O and the following structural formula: H C2H6O H H C C H H Chemical formula O H Structural formula Molecular model Ethyl alcohol A structural formula uses lines between atoms to indicate the covalent bonds Thus, the carbon atoms in ethyl alcohol are covalently bonded to each other, the oxygen atom is bonded to one of the carbon atoms, and the hydrogen atoms are distributed to one carbon, to the other carbon, and to the oxygen Structural formulas are particularly important in organic chemistry—the chemistry of carbon compounds—where the behavior of large, complex molecules is almost entirely governed by their structure Take even a relatively simple substance like glucose, for instance The molecular formula of glucose, C6H12O6, tells nothing about how the atoms are connected In fact, you could probably imagine a great many different ways in which the 24 atoms might be connected The structural formula for glucose, however, shows that carbons and oxygen form a ring of atoms, with the remaining oxygens each bonded to hydrogen and bonded to different carbons H O C H C H C H H O O O O H H H C C H O H C H H [Red = O, gray = C, ivory = H] Glucose—C6H12O6 Some elements even exist as molecules rather than as individual atoms Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine all exist as diatomic (twoatom) molecules whose two atoms are held together by covalent bonds We therefore have to write them as such—H2, N2, O2, F2, Cl2, Br2, and I2—when using any of these elements in a chemical equation Notice that all these diatomic elements except hydrogen cluster toward the far right side of the periodic table 1A 8A H2 2A 3A 4A 5A 6A 7A N2 O2 F2 3B 4B 5B 6B 7B 8B 1B 2B Cl2 Br2 I2 2.12 Ions and Ionic Bonds 61 CONCEPTUAL WORKED EXAMPLE 2.9 Converting between Structural and Molecular Formulas Propane, C3H8, has a structure in which the three carbon atoms are bonded in a row, each end carbon is bonded to three hydrogens, and the middle carbon is bonded to two hydrogens Draw the structural formula, using lines between atoms to represent covalent bonds ▶▶CONCEPTUAL APPLY 2.18 Adrenaline, the so-called flightor-fight hormone, can be represented by the following balland-stick model What is the chemical formula of adrenaline? (Gray = C, ivory = H, red = O, blue = N) SOLUTION H H H H C C C H H H H Propane ▶▶CONCEPTUAL PRACTICE 2.17 Thymine, one of the four bases in deoxyribonucleic acid (DNA), has the following structure What is the chemical formula of thymine? In writing the formula, list the element symbols in alphabetical order, and give the number of each element as a subscript H H H H C N C C C H C N O H O 2.12 IONS AND IONIC BONDS In contrast to a covalent bond, an ionic bond results not from a sharing of electrons but from a transfer of one or more electrons from one atom to another As noted previously, ionic bonds generally form between a metal and a nonmetal Metals, such as sodium, magnesium, and zinc, tend to give up electrons, whereas nonmetals, such as oxygen, nitrogen, and chlorine, tend to accept electrons For example, when sodium metal comes in contact with chlorine gas, a sodium atom gives an electron to a chlorine atom, resulting in the formation of two charged particles, called ions Because a sodium atom loses one electron, it loses one negative charge and becomes an Na+ ion with a charge of +1 Such positive ions are called cations (pronounced cat-ions) Conversely, because a chlorine atom gains an electron, it gains a negative charge and becomes a Cl- ion with a charge of -1 Such negative ions are called anions (an-ions) A sodium atom A sodium cation Na + Cl2 Na+ + ClA chloride anion A chlorine molecule Showing the ions in the preceding chemical equation is useful for keeping track of charged species, but according to convention, ions are not shown in the chemical equation The reaction of sodium and chlorine is written as: Na + Cl ¡ NaCl 2 ▲▲ Chlorine is a toxic green gas, sodium is a reactive metal, and sodium chloride is a harmless white solid 62 chapter Atoms, Molecules, and Ions A similar reaction takes place when magnesium and chlorine molecules (Cl2) come in contact to form MgCl2 A magnesium atom transfers an electron to each of two chlorine atoms, yielding the doubly charged Mg2+ cation and two Cl- anions Mg + Cl2 ¡ Mg2+ + Cl- + ClThe chemical equation for reaction of magnesium with chlorine is: Mg + Cl2 ¡ MgCl2 Because opposite charges attract, positively charged cations such as Na+ and Mg experience a strong electrical attraction to negatively charged anions like Cl- , an attraction that we call an ionic bond Unlike what happens when covalent bonds are formed, though, we can’t really talk about discrete Na+ Cl- molecules under normal conditions We can speak only of an ionic solid, in which equal numbers of Na+ and Cl- ions are packed together in a regular way (FIGURE 2.15) In a crystal of table salt, for instance, each Na+ ion is surrounded by six nearby Cl- ions, and each Cl- ion is surrounded by six nearby Na+ ions, but we can’t specify what pairs of ions “belong” to each other as we can with atoms in covalent molecules 2+ ▶▶FIGURE 2.15 The arrangement of Na ∙ and Cl ∙ ions in a crystal of sodium chloride. There is no discrete “molecule” of NaCl Instead, the entire crystal is an ionic solid Na Cl Na+ Cl– ▶▶Figure It Out Which element loses electrons and which element gains electrons when a metal and nonmetal form an ionic compound? In the sodium chloride crystal, each Na+ ion is surrounded by six nearestneighbor Cl– ions … … and each Cl– ion is surrounded by six nearest-neighbor Na+ ions Charged, covalently bonded groups of atoms, called polyatomic ions, are also common—ammonium ion (NH4+ ), hydroxide ion (OH- ), nitrate ion (NO3- ), and the doubly charged sulfate ion (SO42- ) are examples (FIGURE 2.16) You can think of these polyatomic ions as charged molecules because they consist of specific numbers and kinds of atoms joined together by covalent bonds, with the overall unit having a positive or negative charge When writing the formulas of substances that contain more than one of these ions, parentheses are placed around the entire polyatomic unit The formula Ba(NO3)2, for instance, indicates a substance made of Ba2+ cations and NO3- polyatomic anions in a 1:2 ratio We’ll learn how to name compounds with these ions in Section 2.13 There is no net charge on any ionic compound so the total number of positive charges must equal the total number of negative charges Creating an overall neutral charge serves as a guide for writing the formulas of ionic compounds For example, the formula for the ionic compound formed from sodium (Na+ ) and bromide (Br - ) ions is NaBr because one positive charge from the sodium ion balances one negative charge from the bromide ion The formula for the ionic compound formed from sodium (Na+ ) and sulfate (SO42- ) ions is Na2SO4 because two sodium ions each with a +1 charge are needed to balance the -2 charge on sulfate The formula for the ionic compound formed from sodium (Na+ ) and phosphate (PO43- ) ions is Na3PO4 because three sodium ions each with a +1 charge are needed to balance the -3 charge on phosphate Answer: The metal loses electrons to form a positively charged cation The nonmetal gains electrons to form a negatively charged anion 2.13 Naming Chemical Compounds 63 WORKED EXAMPLE 2.10 +1 charge Identifying Ionic and Molecular Compounds + Which of the following compounds would you expect to be ionic and which molecular? A molecular compound has covalent bonds (a) BaF2 (b) SF4 (c) PH3 (d) CH3OH STRATEGY Remember that covalent bonds generally form between nonmetal atoms, while ionic bonds form between metal and nonmetal atoms Ammonium ion NH4+ –1 charge SOLUTION Compound (a) is composed of a metal (barium) and a nonmetal (fluorine) and is likely to be ionic Compounds (b)–(d) are composed entirely of nonmetals and therefore are probably molecular ▶▶PRACTICE 2.19 Which of the following is an ionic compound? (a) LiBr (b) SiCl4 (c) NF3 - Hydroxide ion OH - –1 charge - ▶▶CONCEPTUAL APPLY 2.20 Which of the following drawings most likely represents an ionic compound and which a molecular compound? Explain (a) (b) Nitrate ion NO3- –2 charge 2– 2.13 NAMING CHEMICAL COMPOUNDS In the early days of chemistry, when few pure substances were known, newly discovered compounds were often given fanciful names—morphine, quicklime, potash, and barbituric acid (said to be named by its discoverer in honor of his friend Barbara) to cite a few Today, with more than 40 million pure compounds known, there would be chaos without a systematic method for naming compounds Every chemical compound must be given a name that not only defines it uniquely but also allows chemists (and computers) to know its chemical structure Different kinds of compounds are named by different rules Ordinary table salt, for instance, is named sodium chloride because of its formula NaCl, but common table sugar (C12H22O11) is named b-d-fructofuranosyl-a-d-glucopyranoside because of special rules for carbohydrates (Organic compounds often have quite complex structures and correspondingly complex names, though we’ll not discuss them in this text.) We’ll begin by seeing how to name simple ionic compounds and then introduce additional rules in later chapters as the need arises Sulfate ion SO42- ▲▲FIGURE 2.16 Molecular models of some polyatomic ions Naming Binary Ionic Compounds Binary ionic compounds—those made of only two elements—are named by identifying first the positive ion and then the negative ion The positive ion takes the same name as the element, while the negative ion takes the first part of its name from the element and then adds the ending -ide For example, KBr is named potassium bromide: potassium for the K+ ion and bromide for the negative Br - ion derived from the element bromine LiF CaBr2 AlCl3 Lithium fluoride Calcium bromide Aluminum chloride ▲▲ Morphine, a pain-killing agent found in the opium poppy, was named after Morpheus, the Greek god of dreams 64 chapter Atoms, Molecules, and Ions ▲▲ Crystals of iron(II) chloride tetrahydrate are greenish, and crystals of iron(III) chloride hexahydrate are brownish yellow ▶▶FIGURE 2.17 Main-group cations (blue) and anions (red). A cation bears the same name as the element it is derived from; an anion name has an -ide ending ▶▶Figure It Out What charge is formed on group 2A and 6A elements? Answer: Group 2A = + and group 6A = - ▶▶FIGURE 2.18 Common transition metal ions. Only ions that exist in aqueous solution are shown FIGURE 2.17 shows some common main-group ions, and FIGURE 2.18 shows some common transition metal ions There are several interesting points about Figure 2.17 Notice, for instance, that metals tend to form cations and nonmetals tend to form anions Also note that elements within a given group of the periodic table form ions with the same charge and that the charge is related to the group number Main-group metals usually form cations whose charge is equal to the group number For example, group 1A elements form singly positive ions (M+ , where M is a metal), group 2A elements form doubly positive ions (M2+ ), and group 3A elements form triply positive ions (M3+ ) Main-group nonmetals usually form anions whose charge is equal to the group number in the U.S system minus eight Thus, group 6A elements form doubly negative ions (6 - = -2), group 7A elements form singly negative ions (7 - = -1), and group 8A elements form no ions at all (8 - = 0) We’ll see the reason for this behavior in Chapter Notice also, in both Figures 2.17 and 2.18, that some metals form more than one kind of cation Iron, for instance, forms both the doubly charged Fe2+ ion and the triply charged Fe3+ ion In naming these ions, we distinguish between them by using a Roman numeral in parentheses to indicate the number of charges Thus, FeCl2 is named iron(II) chloride and FeCl3 is iron(III) chloride Alternatively, an older method distinguishes between the ions by using the Latin name of the element (ferrum in the case of iron) together with the ending -ous for the ion with lower charge and -ic for the ion with higher charge Thus, FeCl2 is sometimes called ferrous chloride and FeCl3 is called ferric chloride Although still in use, this older naming system is being phased out, and we’ll rarely use it in this book 18 8A 1A H+ HHydride 2A Li+ Be 2+ Na+ Mg 2+ Al 3+ S 2ClSulfide Chloride K+ Ca 2+ Ga3+ Se 2BrSelenide Bromide Rb+ Sr 2+ In3+ Sn 2+ Sn 4+ Cs+ Ba 2+ Tl+ Tl3+ Pb 2+ Pb 4+ 13 3A 15 5A N3Nitride 3B 4B 5B 6B 7B Sc 3+ Ti3+ V2+ V3+ Cr2+ Cr3+ Mn2+ ▶▶Figure It Out Can the charge on a transition metal ion be predicted from its group number? 14 4A Y3+ 16 6A 17 7A O 2FOxide Fluoride Te 2ITelluride Iodide 8B 10 11 1B 12 2B Fe2+ Fe3+ Co2+ Ni2+ Cu+ Cu2+ Zn2+ Ru3+ Rh3+ Pd2+ Ag+ Cd2+ Hg2+ (Hg2)2+ Answer: No, many transition metals form different charge states The charge on the transition metal ion is specified in the formula or name 2.13 Naming Chemical Compounds 65 Fe2+ Fe3+ Sn2+ Sn4+ Iron(II) ion Iron(III) ion Tin(II) ion Tin(IV) ion Ferrous ion Ferric ion Stannous ion Stannic ion (From the Latin ferrum = iron) (From the Latin stannum = tin) In any neutral compound, the total number of positive charges must equal the total number of negative charges Thus, you can always figure out the number of positive charges on a metal cation by counting the number of negative charges on the associated anion(s) In FeCl2, for example, the iron ion must be Fe(II) because there are two Clions associated with it Similarly, in TiCl3, the titanium ion is Ti(III) because there are three Cl- anions associated with it As a general rule, a Roman numeral is needed for transition-metal compounds to avoid ambiguity In addition, the main-group metals tin (Sn), thallium (Tl), and lead (Pb) can form more than one kind of ion and need Roman numerals for naming their compounds Metals in group 1A and group 2A form only one cation, however, so Roman numerals are not needed WORKED EXAMPLE 2.11 Converting between Names and Formulas for Binary Ionic Compounds Give systematic names for the following compounds: (a) BaCl2 (b) CrCl3 (c) PbS ▶▶PRACTICE 2.21 Write formulas for the following compounds: STRATEGY Name the cation with the name of the element and the anion using the first part of the element name + “ide.” If the cation is a transition metal, then the charge is specified with Roman numerals Figure out the number of positive charges on each transition metal cation by counting the number of negative charges on the associated anion(s) Refer to Figures 2.17 and 2.18 as necessary SOLUTION (a) Barium chloride (d) Fe 2O3 (a) Magnesium fluoride (b) Tin(IV) oxide ▶▶CONCEPTUAL APPLY 2.22 Three binary ionic compounds are represented on the following periodic table: red with red, green with green, and blue with blue Name each, and write its likely formula No Roman numeral is necessary because barium, a group 2A element, forms only Ba2+ (b) Chromium(III) The Roman numeral III is necessary to chloride specify the + charge on chromium (a transition metal) (c) Lead(II) sulfide The sulfide anion (S2-) has a double negative charge, so the lead cation must be doubly positive (d) Iron(III) oxide The three oxide anions (O2-) have a total negative charge of - 6, so the two iron cations must have a total charge of + Thus, each is Fe(III) Naming Compounds with Polyatomic Ions Ionic compounds that contain polyatomic ions are named in the same way as binary ionic compounds: First the cation is identified and then the anion For example, Ba(NO3)2 is called barium nitrate because Ba2+ is the cation and the NO3- polyatomic anion has the name nitrate Unfortunately, there is no simple systematic way of naming the polyatomic ions themselves, so it’s necessary to memorize the names, formulas, and charges of the most common ones, listed in TABLE 2.5 The ammonium ion (NH4+ ) is the only cation on the list; all the others are anions 66 chapter Atoms, Molecules, and Ions TABLE 2.5 Some Common Polyatomic Ions Formula Name Formula Ammonium Singly charged anions (continued) NO2Nitrite Nitrate NO3- Cation NH4+ Singly charged anions Doubly charged anions CH3CO2- Acetate CN- Cyanide ClO- Hypochlorite ClO2- Chlorite ClO3- Chlorate ClO4- Perchlorate H2PO4- Dihydrogen phosphate HCO3- Hydrogen carbonate (or bicarbonate) HSO4OHMnO4 - Name CO32- Carbonate Chromate Hydrogen sulfate (or bisulfate) CrO42Cr2O72O22HPO42SO32SO42S2O32- Hydroxide Triply charged anion Permanganate PO43- Dichromate Peroxide Hydrogen phosphate Sulfite Sulfate Thiosulfate Phosphate Several points about the ions in Table 2.5 need special mention First, note that the names of most polyatomic anions end in -ite or -ate Only hydroxide (OH- ), cyanide (CN- ), and peroxide (O22- ) have the -ide ending Second, note that several of the ions form a series of oxoanions, binary polyatomic anions in which an atom of a given element is combined with different numbers of oxygen atoms—hypochlorite (ClO- ), chlorite (ClO2- ), chlorate (ClO3- ), and perchlorate (ClO4- ), for example When there are only two oxoanions in a series, as with sulfite (SO32- ) and sulfate (SO42- ), the ion with fewer oxygens takes the -ite ending and the ion with more oxygens takes the -ate ending SO32NO2- Sulfite ion (fewer oxygens) Nitrite ion (fewer oxygens) SO42NO3- Sulfate ion (more oxygens) Nitrate ion (more oxygens) When there are more than two oxoanions in a series, the prefix hypo- (meaning “less than”) is used for the ion with the fewest oxygens, and the prefix per- (meaning “more than”) is used for the ion with the most oxygens ClOClO2 - ClO3ClO4 - Hypochlorite ion (less oxygen than chlorite) Chlorite ion Chlorate ion Perchlorate iron (more oxygen than chlorate) Third, note that several pairs of ions are related by the presence or absence of a hydrogen ion The hydrogen carbonate anion (HCO3- ) differs from the carbonate anion (CO32- ) by the presence of H+ , and the hydrogen sulfate anion (HSO4- ) differs from the sulfate anion (SO42- ) by the presence of H+ The ion that has the additional hydrogen is sometimes referred to using the prefix bi-, although this usage is now discouraged; for example, NaHCO3 is sometimes called sodium bicarbonate HCO3HSO4- Hydrogen carbonate (bicarbonate) ion Hydrogen sulfate (bisulfate) ion CO32SO42- Carbonate ion Sulfate ion 2.13 Naming Chemical Compounds 67 WORKED EXAMPLE 2.12 Converting between Names and Formulas for Compounds with Polyatomic Ions Give systematic names for the following compounds: (b) KHSO4 (c) CuCO3 (a) LiNO3 (d) Fe(ClO4)3 ▶▶PRACTICE 2.23 Write the formula for iron(III) carbonate STRATEGY Name the cation first and the anion second Unfortunately, there is no alternative: The names and charges of the common polyatomic ions must be memorized Refer to Table 2.5 if you need help SOLUTION (a) Lithium nitrate Lithium (group 1A) forms only the Li+ ion and does not need a Roman numeral ▶▶CONCEPTUAL APPLY 2.24 The following drawings are those of solid ionic compounds, with red spheres representing the cations and blue spheres representing the anions in each (1) (2) (b) Potassium Potassium (group 1A) forms only the hydrogen sulfate K+ ion (c) Copper(II) carbonate The carbonate ion has a - charge, so copper must be + A Roman numeral is needed because copper, a transition metal, can form more than one ion (d) Iron(III) perchlorate There are three perchlorate ions, each with a - charge, so the iron must have a + charge Which of the following formulas are consistent with each drawing? (a) LiBr (b) NaNO2 (c) CaCl2 Naming Binary Molecular Compounds Binary molecular compounds—those made of only two covalently bonded elements— are named in much the same way as binary ionic compounds One of the elements in the compound is more electron-poor, or cationlike, and the other element is more electron-rich, or anionlike As with ionic compounds, the cationlike element takes the name of the element itself, and the anionlike element takes an -ide ending The compound HF, for example, is called hydrogen fluoride HF Hydrogen is more cationlike because it is farther left in the periodic table, and fluoride is more anionlike because it is farther right The compound is therefore named hydrogen fluoride We’ll see a quantitative way to decide which element is more cationlike and which is more anionlike in Section 7.3 but you might note for now that it’s usually possible to decide by looking at the relative positions of the elements in the periodic table The farther left and toward the bottom of the periodic table an element occurs, the more likely it is to be cationlike; the farther right and toward the top an element occurs (except for the noble gases), the more likely it is to be anionlike More anionlike More cationlike 68 chapter Atoms, Molecules, and Ions TABLE 2.6 Numerical Prefixes for Naming Compounds Prefix Meaning mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- 10 The following examples show how this generalization applies: CO Carbon monoxide (C is in group 4A; O is in group 6A) CO2 Carbon dioxide PCl3 Phosphorus trichloride (P is in group 5A; Cl is in group 7A) SF4 Sulfur tetrafluoride (S is in group 6A; F is in group 7A) N2O4 Dinitrogen tetroxide (N is in group 5A; O is in group 6A) Because nonmetals often combine with one another in different proportions to form different compounds, numerical prefixes are usually included in the names of binary molecular compounds to specify the numbers of each kind of atom present The compound CO, for example, is called carbon monoxide, and CO2 is called carbon dioxide TABLE 2.6 lists the most common numerical prefixes Note that when the prefix ends in a or o (but not i) and the anion name begins with a vowel (oxide, for instance), the a or o on the prefix is dropped to avoid having two vowels together in the name Thus, we write carbon monoxide rather than carbon monooxide for CO and dinitrogen tetroxide rather than dinitrogen tetraoxide for N2O4 A mono- prefix is not used for the atom named first: CO2 is called carbon dioxide rather than monocarbon dioxide WORKED EXAMPLE 2.13 Converting between Names and Formulas for Binary Molecular Compounds Give systematic names for the following compounds: (b) N2O3 (c) P4O7 (a) PCl3 (d) BrF3 STRATEGY Look at a periodic table to see which element in each compound is more cationlike (located farther to the left or lower) and which is more anionlike (located farther to the right or higher) Then name the compound using the appropriate numerical prefix to specify the number of atoms SOLUTION (a) Phosphorus trichloride (c) Tetraphosphorus heptoxide (b) Dinitrogen trioxide (d) Bromine trifluoride ▶▶PRACTICE 2.25 Write the formula for dinitrogen pentoxide ▶▶CONCEPTUAL APPLY 2.26 Give systematic names for the following compounds: (a) Purple = P, green = Cl (b) Blue = N, red = O ● How Can Measurements of Oxygen and Hydrogen Isotopes in Ice Cores Determine Past Climates? 69 How can measurements of oxygen and hydrogen isotopes in ice cores determine past climates? C limate change refers to variations in average weather conditions, temperatures, and rainfall over an extended period of time Climate research also includes tracking the number and severity of extreme weather events such as heat waves, tornadoes, and hurricanes The Earth’s climate has varied over geological time due to a number of natural causes including variations in the Earth’s orbit, the sun’s intensity, particulates from volcanic eruptions, and levels of greenhouse gases The term climate change or global warming is most often associated with the pronounced warming of the climate from the mid- to late twentieth century, which is largely attributed to increased levels of carbon dioxide in the atmosphere from burning fossil fuels Examine FIGURE 2.19, which shows the change in global surface temperature (°C) from 1880 to the present relative to the long-term average temperature from 1901–2000 Temperatures measured on land and at sea show that Earth’s globally averaged surface temperature is rising Though warming has not been uniform across the planet, the upward trend in the globally averaged temperature shows that more areas are warming 0 1880 1900 1920 1940 1960 1980 Anomaly (°F) Anomaly (°C) than cooling For the past 45 years, global surface temperature rose at an average rate of about 0.17 °C (about 0.3 °Fahrenheit) per decade—more than twice as fast as the 0.07 °C per decade increase observed for the entire period of recorded observations (1880–2015) Remarkably, all 16 years of the twenty-first century rank among the 17 warmest years on record In recent history, scientists have been recording the Earth’s average temperature using satellite measurements and data from numerous weather and research stations Surrounding temperatures and other information are used to fill in data from areas that have few measurements This process provides a consistent, reliable method for monitoring changes in Earth’s surface temperature over time However, a long-term record of past climate is needed to put the recent warming trends in context Analyzing a historical temperature record enables scientists to evaluate previous rates of temperature change and the magnitude of temperature changes that caused drastic differences in climate such as ice ages Clues to past climates are etched on our planet in polar ice caps, cave rocks, coral reefs, and tree rings Measuring oxygen and hydrogen isotope ratios in polar and glacial ice can be used to reconstruct past global temperatures Sensitive mass spectrometers (Section 2.10) are used to measure isotope ratios in water from ice core samples FIGURE 2.20 shows a near linear relationship between the difference in the 18O/16O ratio in snowfall and the mean annual temperature for that site Isotopic ratios are a measure of temperature because more energy Change in 18O (percent) INQUIRY -2 -4 2000 ▲▲FIGURE 2.19 Global land and ocean temperature anomalies Annual global temperatures since 1880 compared to the longterm average temperature from 1901–2000 The zero line represents the long-term average global temperature, and the blue and red bars show the difference above or below average for each year Data come from a combined set of land-based weather stations and sea-surface temperature measurements -6 -60 -40 -20 20 Temperature (°C) ▲▲FIGURE 2.20 Difference in the 18O> 16O ratio in ice core samples related to the average temperature when the snowfall occurred 70 chapter Atoms, Molecules, and Ions Temperature difference Change in deuterium levels (percent) is required to evaporate water molecules containing a heavy isotope from the surface of the ocean than water molecules with lighter isotopes As warm air is transported to cold, polar regions, water molecules containing heavier isotopes preferentially precipitate Therefore, the ratio of heavier isotopes to lighter isotopes (18O/16O and 2H/1H) in precipitation increases with warmer temperatures Cold locations such as Antarctica have about 5% less 18O than warm ocean water Plotting either 18O/16O or 2H/1H with ice core depth reveals oscillations in temperature as a function of time Ice core samples are dated by number of layers and depth The data in FIGURE 2.21 was generated from an Antarctic ice core which extends km in length and dates back 800,000 years The top graph shows variations in the amount of the heavy isotope of water (2H, deuterium), and the bottom graph shows the correlation with temperature The data reveals cold glacial periods interspersed by warm periods roughly every 100,000 years Historical climate records show that the Earth has experienced warm and cool periods, but the rate of warming in the past 50 years is unprecedented Past global temperatures can also be correlated to levels of greenhouse gases such as carbon dioxide trapped in bubbles in the ice Section 10.11 discusses the effect of greenhouse gases on climate, and Figure 10.25 shows the correlation between carbon dioxide concentration and global temperature -360 -380 -400 -420 -440 PROBLEM 2.28 Use Figure 2.19 to determine if the following statements are true or false (a) Each year after 1950 had an average global temperature higher than the 1901–2000 average temperature (b) Prior to 1940, average global temperatures were lower than the 1901–2000 average temperature (c) Warming in the time period 1970–2015 occurred at a faster rate than warming in the time period 1900–1950 (d) Global average temperatures exhibited a cooling trend from 1880–1910 PROBLEM 2.29 How many protons, neutrons, and electrons are in 18O and 2H? PROBLEM 2.30 Which sample of H2O has a higher ratio of 18 O/16O: warm seawater near the equator or snow falling in Antarctica? PROBLEM 2.31 The last ice age occurred from 110,000 to 11,700 years ago Use Figure 2.21 to answer questions about variations in temperature and change in deuterium percent between a warm climate and an ice age (a) What is the difference in temperature from our current warm climate to the maximum extent of glaciation occurring approximately 22,000 years ago? (b) What is the change in deuterium percent during the same time period? PROBLEM 2.32 For this problem, assume that water consists only of the most abundant isotopes of oxygen (16O and 18O) The atomic mass for 16O is 15.9949146 u, and the atomic mass for 18 O is 17.9991610 u (a) A standard seawater sample contains 0.1995% 18O Calculate the atomic weight of oxygen in seawater, and report your answer to five decimal places (b) A polar ice core sample contains 0.1971% 18O Calculate the atomic weight of oxygen in polar ice, and report your answer to five decimal places -460 -480 -2 -4 -6 -8 -10 -12 800,000 PROBLEM 2.27 Global climate is affected by variations in (a) the Earth’s orbit around the Sun (b) particulates in the atmosphere from pollution or volcanoes (c) the Sun’s intensity (d) greenhouse gas levels (e) all of the above 600,000 400,000 200,000 Age (years before present) ▲▲FIGURE 2.21 EPICA Dome C ice core 800 kYr deuterium data and temperature estimate PROBLEM 2.33 For isotopic analysis, an ice core sample was heated to produce gaseous H2O If 1.00 mg of gaseous H2O was injected into a mass spectrometer: (a) How many moles of water were injected? (b) If the sample contains 0.0156% deuterium, how many deuterium atoms were injected? Study Guide 71 READY-TO-GO STUDY TOOLS in the Mastering Chemistry Study Area help you master the toughest topics in General Chemistry Problem-Solving videos and Practice Tests are all in one, easy to navigate place to help keep you focused and give you the support you need to succeed STUDY GUIDE Section Concept Summary Learning Objectives Test Your Understanding 2.1 Chemistry and the Elements All matter is formed from one or more of the 118 presently known elements—fundamental substances that can’t be chemically broken down Elements are symbolized by one- or two-letter abbreviations 2.1 Write symbols to represent element names Problems 2.48, 2.50, 2.52 2.2 Elements and the Periodic Table Elements are organized into a periodic table with groups (columns) and periods (rows) Elements in the same groups show similar chemical behavior Elements are classified as metals, nonmetals, or semimetals 2.2 Identify the location of metals, nonmetals, and semimetals on the periodic table Problems 2.37, 2.62, 2.63 2.3 Indicate the atomic number, group number, and period number for an element whose position in the periodic is given Problem 2.36 2.4 Identify groups as main group, transition, metal group, or inner transition metal group Problems 2.59, 2.61 The characteristics, or properties, that are used to describe matter can be classified in several ways Physical properties are those that can be determined without changing the chemical composition of the sample, whereas chemical properties are those that involve a chemical change Intensive properties are those whose values not depend on the size of the sample, whereas extensive properties are those that depend on sample amount 2.5 Specify the location and give examples of elements in the alkali metal, alkaline earth metal, halogen, noble gas groups Problem 2.35, 2.64–2.67 2.6 Classify an element as a metal, nonmetal, or semimetal using its properties Worked Example 2.1; Problems 2.1–2.2, 2.68–2.71 2.4 Observations Supporting Atomic Theory: The Conservation of Mass and the Law of Definite Proportions Elements join together in different ways to make chemical compounds, and a pure compound always has the same proportion of elements by mass During a chemical reaction, the law of mass conservation applies, and the mass of reactants is the same as the mass of products 2.7 Determine the mass of the products in a reaction using the law of mass conservation Problems 2.78–2.79 2.5 The Law of Multiple Proportions and Dalton’s Atomic Theory Elements are made of tiny particles called atoms, which can combine in simple numerical ratios according to the law of multiple proportions 2.8 Demonstrate the law of multiple proportions using mass composition of two compounds of the same elements Worked Example 2.2; Problems 2.80, 2.82 2.9 Determine the formula of a compound given mass composition data for two compounds and the formula of one compound Problems 2.83–2.84 2.6 Atomic Structure: Electrons Atoms are composed of three fundamental particles: protons are positively charged, electrons are negatively charged, and neutrons are neutral 2.10 Describe Thomson’s cathode-ray experiment and what it contributed to the current model of atomic structure Problems 2.86–2.88 2.11 Describe Millikan’s oil drop experiment and what it contributed to the current model of atomic structure Problems 2.89–2.90 2.12 Describe Rutherford’s gold foil experiment and what it contributed to the current model of atomic structure Problems 2.91–2.92 2.13 Calculate the number of atoms in sample given the size of the atom Problems 2.93–2.94 2.3 Some Common Groups of Elements and Their Properties 2.7 Atomic Structure: Proton and Neutrons According to the nuclear model of an atom proposed by Ernest Rutherford, protons and neutrons are clustered into a dense core called the nucleus, while electrons move around the nucleus at a relatively great distance 72 chapter Atoms, Molecules, and Ions Section Concept Summary Learning Objectives Test Your Understanding 2.8 Atomic Numbers Elements differ from one another according to how many protons their atoms contain, a value called the atomic number (Z) of the element The sum of an atom’s protons and neutrons is its mass number (A) Although all atoms of a specific element have the same atomic number, different atoms of an element can have different mass numbers depending on how many neutrons they have Atoms with identical atomic numbers but different mass numbers are called isotopes 2.14 Determine the mass number, atomic number, and number of protons neutrons and electrons from an isotope symbol Worked Example 2.4; Problems 2.100–2.102, 2.102, 2.104, 2.106 2.9 Atomic Weights and the Mole Atomic weights are measured using the unified atomic mass unit (u), defined as 1/12 the mass of a 12C atom Because both protons and neutrons have a mass of approximately 1, the mass of an atom in unified atomic mass units is numerically close to the atom’s mass number The element’s atomic weight is a weighted average of the isotopic masses of its naturally occurring isotopes When referring to the enormous numbers of atoms that make up visible amounts of matter, the fundamental SI unit called a mole is used One mole is the amount whose mass in grams, called its molar mass, is numerically equal to the atomic weight Numerically, one mole of any element contains 6.022 * 1023 atoms, a value called Avogadro’s number (NA) 2.15 Calculate atomic weight given the fractional abundance and mass of each isotope Worked Example 2.5; Problems 2.116, 2.118, 2.120 2.16 Convert between grams and numbers of moles or atoms using molar mass and Avogadro’s number Worked Example 2.6; Problems 2.124–2.125 2.10 Measuring Atomic Weight: Mass Spectrometry A mass spectrometer separates gaseous ions based on their mass-to-charge ratio The mass spectrum records the intensity (the number) of ions on the y-axis and the mass-to-charge ratio of the ions on the x-axis Mass spectral data can be used to calculate the atomic weight of an element 2.17 Use data from a mass spectrum to calculate the atomic weight of an element Worked Example 2.7; Problems 2.13–2.14, 2.132–2.133 2.11 Mixtures and Chemical Compounds; Molecules and Covalent Bonds Most substances are chemical compounds, formed when atoms of two or more elements combine in a chemical reaction The atoms in a compound are held together by one of two kinds of chemical bonds Covalent bonds form when two atoms share electrons to give a new unit of matter called a molecule 2.18 Classify matter as a mixture, pure substance, element, or compound Worked Example 2.8; Problems 2.15–2.16 2.42 2.19 Convert between structural formulas, ball-and-stick models, and chemical formulas Worked Example 2.9; Problems 2.43–2.44 2.12 Ions and Ionic Bonds Ionic bonds form when one atom completely transfers one or more electrons to another atom, resulting in the formation of ions Positively charged ions (cations) are strongly attracted to negatively charged ions (anions) by electrical forces 2.20 Classify bonds as ionic or covalent Worked Example 2.10; Problems 2.134–2.135 2.21 Determine the number of electrons and protons from chemical symbol and charge Problems 2.41, 2.138–2.139 2.22 Match the molecular representation of an ionic compound with its chemical formula Problem 2.45 2.23 Convert between name and formula for ionic compounds Worked Examples 2.11–2.12; Problems 2.146, 2.148, 2.150, 2.152 2.24 Convert between name and formula for binary molecular compounds Worked Example 2.13; Problems 2.161–2.162 2.13 Naming Chemical Compounds Chemical compounds are named systematically by following a series of rules Binary ionic compounds are named by identifying first the positive ion and then the negative ion Binary molecular compounds are similarly named by identifying the cationlike and anionlike elements Naming compounds with polyatomic ions involves memorizing the names and formulas of the most common ones Practice Test 73 KEY TERMS covalent bond 58 electron 45 element 34 extensive property 38 group 37 inner transition metal group 38 intensive property 38 ion 61 ionic bond 61 ionic solid 62 isotope 50 anion 61 atom 44 atomic mass 52 atomic number (Z) 49 atomic weight 52 Avogadro’s number 53 cation 61 chemical bond 58 chemical compound 42 chemical equation 42 chemical formula 42 chemical property 38 law of definite proportions 43 law of mass conservation 42 law of multiple proportions 43 main group 37 mass number 50 mass spectrometer 55 mass spectrum 55 mixture 57 molar mass 53 mole 53 molecule 58 neutron 48 nucleus 47 oxoanion 66 period 37 periodic table 35 physical property 38 polyatomic ion 62 property 38 proton 48 structural formula 60 transition metal group 38 unified atomic mass unit (u) 52 PRACTICE TEST After studying this chapter, you can assess your understanding with these practice test questions, which are correlated with chapter learning objectives If you answer a question incorrectly, refer to the learning objectives in the end-of-chapter Study Guide for assistance The Study Guide provides a conceptual summary, references a Worked Example to model how to solve the problem, and gives additional problems for more practice Refer to a periodic table Which pair of elements you expect to be most similar in their chemical properties? (LO 2.3) (a) K and Cu (b) O and Se (c) Be and B (d) Rb and Sr Identify the location of the element in period 4, group 6A and classify it as a metal, nonmetal, or semimetal (LO 2.2) d a b c (a) Element in position a; nonmetal (b) Element in position b; metal (c) Element in position c; semimetal (d) Element in position d; metal Which description of an element is incorrectly matched with its location in the periodic table? (LO 2.5–2.6) (a) Element 3—An element in the transition metal group that is a good conductor of electricity (b) Element 2—An element that is in the halogen group and does not conduct electricity (c) Element 4—An element in alkali metal group that is found in its pure form in nature (d) Element 1—An element that is a solid at room temperature, brittle, and a poor conductor of electricity A compound containing sulfur and fluorine contains 8.00 g of S and 9.50 g of F Which combination of S and F masses represents a different compound that obeys the Law of Multiple Proportions? (LO 2.8) (a) 32.0 g of S and 38.0 g of F (b) 4.00 g of S and 4.75 g of F (c) 8.00 g of S and 10.5 g of F (d) 16.0 g of S and 57.0 g of F Which experiment and subsequent observation led to the discovery that atoms contain negatively charged particles, now known as electrons? (LO 2.10–2.12) (a) Oil is sprayed into a chamber and the speed at which the oil droplets fall is measured with and without an applied voltage X rays in the chamber knock electrons out of air molecules The electrons stick to the oil producing an overall negative charge on the drops Adjusting the voltage changes the speed at which the negatively charged oil droplets fall (b) When a high voltage is applied across metal electrodes at opposite ends of a sealed glass tube, a cathode ray is produced The cathode ray is repelled by a negatively charged plate (c) A radioactive substance emits alpha particles, which are directed at a thin gold foil Most of the alpha particles pass through the foil, but a few alpha particles are slightly deflected and some even bounce back toward the radioactive source (d) The mass of different elements in a pure chemical compound are measured Different samples of the compound always contains the same proportion of elements by mass How many protons, neutrons, and electrons are present in an atom of 206 82 Pb? (LO 2.14) (a) 82 protons, 206 neutrons, 82 electrons (b) 124 protons, 82 neutrons, 124 electrons (c) 82 protons, 124 neutrons, 82 electrons (d) 82 protons, 82 neutrons, 124 electrons ... Data Names: Robinson, Jill K | McMurry, John | Fay, Robert C., 1936Title: Chemistry / Jill K Robinson (Indiana University), John E McMurry (Cornell University), Robert C Fay (Cornell University)... 1A Main groups Periodic Table of the Elements CHEMISTRY E I G H T H JILL K ROBINSON Indiana University JOHN E MCMURRY Cornell University ROBERT C FAY Cornell University E D I T I O N Director... colleagues at so many other institutions who read, criticized, and improved our work Jill K Robinson John McMurry Robert C Fay For Instructors xxiii REVIEWERS FOR THE EIGHTH EDITION Stanley Bajue, Medger
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