https://1drv.ms/b/s!AmkCsf2WlV7n1xM-c5g_7KYRWpZZ?e=5GaAhX
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 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 74 chapter Atoms, Molecules, and Ions What is the atomic weight of an element that consists of two naturally occurring isotopes? The first isotope has a mass of 84.911 and an abundance of 72.17% and the second isotope has a mass of 86.909 and an abundance of 27.83% (LO 2.15) (a) 85.47 (b) 86.35 (c) 85.91 (d) 85.17 Which sample has the greatest mass? (LO 2.16) (a) 5.5 mol of C (b) 2.1 mol of S (c) 4.2 mol of Be (d) 0.52 mol of Ag How many atoms are present in 1.2 g of gold? (LO 2.16) (a) 2.5 * 1021 (b) 1.4 * 1026 (c) 7.2 * 1023 (d) 3.7 * 1021 10 Bromine has two naturally occurring isotopes; 79Br (mass of 78.918) and 81Br (mass of 80.916) If the atomic weight of bromine is 79.904, predict the mass spectrum of a sample of bromine atoms (LO 2.17) (a) (b) 11 The molecular illustration represents (LO 2.18) 12 13 14 80 80 60 60 Intensity Intensity 100 40 20 40 20 77 78 79 80 81 82 83 Mass/charge 77 78 79 80 81 82 83 Mass/charge (c) (d) 100 100 80 80 60 60 Intensity Intensity 15 40 20 20 40 77 78 79 80 81 82 83 Mass/charge 77 78 79 80 81 82 83 Mass/charge (a) a pure element (b) a mixture of two elements (c) a pure compound (d) a mixture of an element and a compound Which of the following compounds would you expect to have covalent bonds? (LO 2.20) (a) Na2O (b) PBr3 (c) CaBr2 (d) MgS + How many protons, neutrons, and electrons are in 107 47Ag ? (LO 2.21) (a) protons = 47, neutrons = 60, electrons = 46 (b) protons = 47, neutrons = 107, electrons = 48 (c) protons = 60, neutrons = 47, electrons = 47 (d) protons = 47, neutrons = 107, electrons = 46 What is the correct formula for sodium phosphate? (LO 2.24) (a) Na3PO4 (b) Na3P (c) NaPO4 (d) Na(PO4)2 Which of the following compounds is incorrectly named? (LO 2.23–2.25) (a) CaO; calcium oxide (b) FeBr2; iron dibromide (c) N2O5; dinitrogen pentoxide (d) CrO3; chromium(VI) oxide Answers: b, a, c, d, b, c, a, b, d, 10 b, 11 c, 12 b, 13 a, 14 a, 15 b Conceptual Problems 75 Mastering Chemistry RAN provides end-of-chapter exercises, feedback-enriched tutorial problems, animations, and interactive activities to encourage problem-solving practice and deeper understanding of key concepts and topics Randomized in Mastering Chemistry CONCEPTUAL PROBLEMS Problems 2.1–2.33 appear within the chapter 2.34 Where on the following outline of a periodic table are the indicated elements or groups of elements? (d) Transition metals (a) Alkali metals (e) Hydrogen (b) Halogens (f) Helium (c) Alkaline earth metals 2.35 Is the red element on the following periodic table likely to be a gas, a liquid, or a solid? What is the atomic number of the blue element? What is the group number of the green, blue, and red elements? Name at least one other element that is chemically similar to the green element 2.38 If yellow spheres represent sulfur atoms and red spheres represent oxygen atoms, which of the following drawings shows a collection of sulfur dioxide (SO2) units? (a) (b) (c) (d) 2.39 Assume that the mixture of substances in drawing (a) undergoes a reaction Which of the drawings (b)–(d) represents a product mixture consistent with the law of mass conservation? (a) 2.36 The element indicated on the following periodic table is used in smoke detectors Identify it, give its atomic number, and tell what kind of group it’s in (b) 2.37 Identify the three elements indicated on the periodic table, RAN and give the group that they are in Classify these elements as metals, nonmetals, or semimetals Would you expect these elements to have similar or different chemical reactivity? 1A 8A 2A 3A 4A 5A 6A 7A (c) (d) 76 chapter Atoms, Molecules, and Ions 2.40 In the following drawings, red spheres represent protons, RAN and blue spheres represent neutrons Which of the drawings represent different isotopes of the same element, and which represents a different element altogether? (a) (b) (a) Alanine (an amino acid) (c) 2.41 Which of the following three drawings represents a neutral Na atom, which represents a Ca atom with two positive electrical charges (Ca2+), and which represents an F atom with one minus charge (F-)? (a) (b) 9+ 20+ 11- 18- 2.42 In the following drawings, red and blue spheres represent atoms of different elements Match the molecular pictures (a)–(c) with the following descriptions: (i) a pure substance consisting of a compound (ii) a pure substance consisting of an element (iii) a mixture of elements (b) (a) (c) 2.43 Methionine, one of the 20 amino acid building blocks from RAN which proteins are made, has the following structure What is the chemical formula of methionine? In writing the formula, list the element symbols in alphabetical order, and give the number of each element as a subscript H H H C S H H C H C C H H O C N H (b) Ethylene glycol (automobile antifreeze) (c) 11+ 10- 2.44 Give molecular formulas corresponding to each of the following ball-and-stick molecular representations (red = O, gray = C, blue = N, ivory = H) In writing the formula, list the elements in alphabetical order O H H Methionine (an amino acid) (c) Acetic acid (vinegar) 2.45 In the following drawings, red spheres represent cations, and blue spheres represent anions Match the drawings (a)–(d) with the following ionic compounds: (i) Ca3(PO4)2 (ii) Li2CO3 (iii) FeCl2 (iv) MgSO4 (a) (b) (c) (d) Section Problems 77 SECTION PROBLEMS Elements and the Periodic Table (Sections 2.1–2.3) 2.46 How many elements are presently known? About how many occur naturally? 2.47 Which element accounts for roughly 75% of the observed mass of the universe? Which four elements make up 95% of the mass of the human body? 2.48 Look at the alphabetical list of elements inside the front RAN cover What are the symbols for the following elements? (a) Gadolinium (used in color TV screens) (b) Germanium (used in semiconductors) (c) Technetium (used in biomedical imaging) (d) Arsenic (used in pesticides) 2.49 Look at the alphabetical list of elements inside the front RAN cover What are the symbols for the following elements? (a) Cadmium (used in rechargeable Ni-cad batteries) (b) Iridium (used for hardening alloys) (c) Beryllium (used in the space shuttle) (d) Tungsten (used in light bulbs) 2.50 Look at the alphabetical list of elements inside the front cover RAN Give the names corresponding to the following symbols: (a) Te (b) Re (c) Be (d) Ar (e) Pu 2.51 Look at the alphabetical list of elements inside the front cover RAN Give the names corresponding to the following symbols: (a) B (b) Rh (c) Cf (d) Os (e) Ga 2.52 What is wrong with each of the following statements? (a) The symbol for tin is Ti (b) The symbol for manganese is Mg (c) The symbol for potassium is Po (d) The symbol for helium is HE 2.53 What is wrong with each of the following statements? (a) The symbol for carbon is ca (b) The symbol for sodium is So (c) The symbol for nitrogen is Ni (d) The symbol for chlorine is Cr 2.54 Examine Figure 2.2, A portion of Mendeleev’s periodic table (a) Which characteristic was used to organize the elements in the table? (b) How was Mendeleev able to predict with a high level of accuracy the properties of the undiscovered elements beneath Al and Si in the periodic table? 2.55 Examine Figure 2.3, The periodic table (a) Which characteristic was used to organize the elements in the table? (b) A horizontal row is called a (c) A vertical column is called a (d) Groups 1A–8A (also known as 1, 2, and 13–18) are called groups (e) Groups 1B–8B (also known as 3–12) are called groups (f) Classify the following three elements as metals, nonmetals, or semimetals: Mo, Br, Si (g) What is the symbol for the element located in period 3, group 4A? 2.56 What are the rows called and what are the columns called in the periodic table? 2.57 How many groups are there in the periodic table? How are they labeled? 2.58 What common characteristics elements within a group of the periodic table have? 2.59 Where in the periodic table are the main-group elements found? Where are the transition metal groups found? 2.60 Where in the periodic table are the metallic elements found? Where are the nonmetallic elements found? 2.61 What is a semimetal, and where in the periodic table are semimetals found? 2.62 Classify the following elements as metals, nonmetals, or RAN semimetals: (a) Ti (b) Te (c) Se (d) Sc (e) Si 2.63 Classify the following elements as metals, nonmetals, or RAN semimetals: (a) Ar (b) Sb (c) Mo (d) Cl (e) N (f) Mg 2.64 List several general properties of the following groups: (a) Alkali metals (b) Noble gases (c) Halogens 2.65 (a) Without looking at a periodic table, list as many alkali RAN metals as you can (There are five common ones.) (b) Without looking at a periodic table, list as many alkaline earth metals as you can (There are five common ones.) 2.66 Without looking at a periodic table, list as many halogens as you can (There are four common ones.) 2.67 Without looking at a periodic table, list as many noble gases as you can (There are six common ones.) 2.68 At room temperature, a certain element is found to be a soft, silver-colored solid that reacts violently with water and is a good conductor of electricity Is the element likely to be a metal, a nonmetal, or a semimetal? 2.69 At room temperature, a certain element is found to be a shiny, silver-colored solid that is a poor conductor of electricity When a sample of the element is hit with a hammer, it shatters Is the element likely to be a metal, a nonmetal, or a semimetal? 2.70 At room temperature, a certain element is yellow crystalline RAN solid It does not conduct electricity and when hit with a hammer, it shatters Is the element likely to be a metal, a nonmetal, or a semimetal? 2.71 At room temperature, a certain element is a colorless, unreRAN active gas Is the element likely to be a metal, a nonmetal, or a semimetal? 2.72 In which of the periodic groups 1A, 2A, 5A, and 7A is the first letter of all elements’ symbol the same as the first letter of their name? 2.73 For which elements in groups 1A, 2A, 5A, and 7A of the periodic table does the first letter of their symbol differ from the first letter of their name? 2.74 Which type of property—intensive or extensive—does not depend on the amount of substance present? 2.75 Label the following properties as intensive or extensive: density, volume, mass, electrical conductivity 78 chapter Atoms, Molecules, and Ions Atomic Theory (Sections 2.4–2.5) Elements and Atoms (Sections 2.6–2.8) 2.76 How does Dalton’s atomic theory account for the law of mass conservation and the law of definite proportions? 2.77 What is the law of multiple proportions, and how does Dalton’s atomic theory account for it? 2.78 A sample of mercury with a mass of 114.0 g was combined RAN with 12.8 g of oxygen gas, and the resulting reaction gave 123.1 g of mercury(II) oxide How much oxygen was left over after the reaction was complete? 2.79 A sample of CaCO3 was heated, causing it to form CaO RAN and CO2 gas Solid CaO remained behind, while the CO2 escaped to the atmosphere If the CaCO3 weighed 612 g and the CaO weighed 343 g, how many grams of CO2 were formed in the reaction? 2.80 In methane, one part hydrogen combines with three parts carbon by mass If a sample of a compound containing only carbon and hydrogen contains 32.0 g of carbon and 8.0 g of hydrogen, could the sample be methane? If the sample is not methane, show that the law of multiple proportions is followed for methane and this other substance 2.81 In borane, one part hydrogen combines with 3.6 parts boron by mass A compound containing only hydrogen and boron contains 6.0 g of hydrogen and 43.2 g of boron Could this compound be borane? If it is not borane, show that the law of multiple proportions is followed for borane and this other substance 2.82 Benzene, ethane, and ethylene are just three of a large number of hydrocarbons—compounds that contain only carbon and hydrogen Show how the following data are consistent with the law of multiple proportions 2.86 The results from Thomson’s cathode-ray tube experiment RAN led to the discovery of which subatomic particle? 2.87 What affects the magnitude of the deflection of the cathode ray in Thomson’s experiment? 2.88 Label the following statements about J J Thomson’s cathoderay tube experiments shown in Figure 2.6 as true or false (a) When a high voltage is applied to metal electrodes in a sealed, evacuated glass tube, an electric current flows (b) A cathode ray is a stream of charged particles (c) The cathode ray is deflected away from a positively charged plate (d) Many different types of metal electrodes are capable of producing a cathode ray (e) A cathode ray is made up of protons (f) By measuring the deflection of the cathode ray beam caused by electric fields of known strength, the chargeto-mass ratio of the electron was calculated 2.89 Fill in the blanks in the description of Millikan’s oil drop experiment, shown in Figure 2.7 A fine mist of oil was sprayed into a chamber, and the velocity of the oil droplets was measured using a telescopic eyepiece Knowing the terminal velocity of a falling oil droplet allows the (mass/charge/volume) of the oil droplet to be calculated Energetic X rays remove (electrons/protons) from air molecules, which adhere to the surface of the oil droplet, giving it a (negative/positive) charge The charged oil droplets fall between two horizontal plates with an applied voltage The top plate is given a (negative/positive) charge, which slows the fall of the drop and suspends it between the two plates The overall charge on the drop can be calculated from the mass of the drop and the voltage on the plates The charge on the oil drop was found to be (a fraction/a whole number multiple) of the charge of an electron The oil drop experiment measured the (mass/charge) of the electron, allowing the (mass/charge) of the electron to be calculated from J J Thomson’s charge-to-mass ratio measurements 2.90 Which of the following charges is not possible for the overRAN all charge on an oil droplet in Millikan’s experiment? For this problem, we’ll round the currently accepted charge of an electron to 1.602 * 10-19 C (a) - 1.010 * 10-18 C (b) - 8.010 * 10-19 C -18 (c) - 2.403 * 10 C 2.91 What discovery about atomic structure was made from the results of Rutherford’s gold foil experiment? 2.92 Prior to Rutherford’s gold foil experiment, the “plum pudding” model of the atom represented atomic structure In this model, the atom is composed of electrons interspersed within a positive cloud of charge If this were the correct model of the atom, predict how the results of Rutherford’s experiment would have been different (a) The alpha particles would pass right through the gold foil with little to no deflection (b) Most of the alpha particles would be deflected back toward the source Mass of Carbon in 5.00 g Sample Mass of Hydrogen in 5.00 g Sample Benzene 4.61 g 0.39 g Ethane 4.00 g 1.00 g Ethylene 4.29 g 0.71 g Compound 2.83 The atomic weight of carbon (12.011) is approximately 12 times that of hydrogen (1.008) (a) Show how you can use this knowledge to calculate possible formulas for benzene, ethane, and ethylene (Problem 2.82) (b) Show how your answer to part (a) is consistent with the actual formulas for benzene (C6H6), ethane (C2H6), and ethylene (C2H4) 2.84 Two compounds containing carbon and oxygen have the RAN following percent composition by mass Compound 1: 42.9% carbon and 57.1% oxygen Compound 2: 27.3% carbon and 72.7% oxygen Show that the law of multiple proportions is followed If the formula of the first compound is CO, what is the formula of the second compound? 2.85 In addition to carbon monoxide (CO) and carbon dioxide RAN (CO2), there is a third compound of carbon and oxygen called carbon suboxide If a 2.500 g sample of carbon suboxide contains 1.32 g of C and 1.18 g of O, show that the law of multiple proportions is followed What is a possible formula for carbon suboxide? Section Problems 79 (c) Most of the alpha particles would be absorbed by the atom and neither pass through nor be deflected from the gold foil 2.93 A period at the end of sentence written with a graphite penRAN cil has a diameter of mm If the period represented the nucleus, approximately how large is the diameter of the entire atom in units of m? 2.94 A 1/4@inch@thick lead sheet is used for protection from mediRAN cal X rays If a single lead atom has a diameter 350 pm, how many atoms thick is the lead sheet? 2.95 A period at the end of sentence written with a graphite penRAN cil has a diameter of mm How many carbon atoms would it take to line up across the period if a single carbon atom has a diameter of 150 pm? 2.96 What is the difference between an atom’s atomic number and its mass number? 2.97 What is the difference between an element’s atomic number and its atomic weight? 2.98 The subscript giving the atomic number of an atom is often left off when writing an isotope symbol For example, 136C is often written simply as 13C Why is this allowed? 2.99 Iodine has a lower atomic mass than tellurium RAN (126.90 for iodine, 127.60 for tellurium) even though it has a higher atomic number (53 for iodine, 52 for tellurium) Explain 2.100 Give the names and symbols for the following elements: (a) An element with atomic number (b) An element with 18 protons in its nucleus (c) An element with 23 electrons 2.101 The radioactive isotope cesium-137 was produced in large amounts in fallout from the 1985 nuclear power plant disaster at Chernobyl, Ukraine Write the symbol for this isotope in standard format 2.102 Write symbols for the following isotopes: RAN (a) Radon-220 (b) Polonium-210 (c) Gold-197 2.103 Write symbols for the following isotopes: (a) Z = 58 and A = 140 (b) Z = 27 and A = 60 2.104 How many protons, neutrons, and electrons are in each of RAN the following atoms? (a) 157N (b) 60 (c) 131 (d) 148 27Co 53I 58Ce 2.105 How many protons and neutrons are in the nucleus of the following atoms? (b) 32S (a) 27Al (c) 64Zn (d) 207Pb 2.106 Identify the following elements: RAN (c) 104 (a) 24 (b) 58 46X 12X 28X 2.107 Identify the following elements: (b) 195 (a) 202 (c) 184 78X 80X 76X (d) 183 74X (d) 209 83X 2.108 Which of the following isotope symbols can’t be correct? 18 12 33 18 11 9F 5C 35Br 8O 5Bo 2.109 Which of the following isotope symbols can’t be correct? RAN 14 131 54 73 7Ni 54Xe 26Fe 23Ge 2He 2.110 Fluorine occurs naturally as a single isotope How many protons, neutrons, and electrons are present in deuterium fluoride (2HF)? (Deuterium is 2H.) 2.111 Hydrogen has three isotopes (1H, 2H, and 3H), and chlorine has two isotopes (35Cl and 37Cl) How many isotopic kinds of HCl are there? Write the formula for each, and tell how many protons, neutrons, and electrons each contains Atomic Weight, Moles, and Mass Spectrometry (Sections 2.9–2.10) 2.112 The unified atomic mass unit (u) is defined as exactly 1/12 the mass of a neutral atom of: (d) 16O (a) 1H (c) 14C (b) 12C 2.113 (a) The unified atomic mass unit (u) is used to represent the extremely small mass of atoms How many grams are equivalent to u? (b) The mole is a unit used to represent a very large number of atoms How many atoms are equivalent to mol of atoms? 2.114 Match the descriptions (a)–(e) with the following terms: atomic weight, atomic mass, mass number, atomic number, molar mass (a) The mass of a specific atom such as one atom of 13C (b) The quantity determined by the number of protons in an element (c) The number of grams in mol of an element (d) The number of protons and neutrons in an element (e) The weighted average of the isotopic masses of an element’s naturally occurring isotopes 2.115 Label the following statements as true or false (a) The atomic weight and the atomic number of element have the same numerical value (b) The molar mass in grams for an element and the atomic weight have the same numerical value 2.116 Copper has two naturally occurring isotopes, including 65 Cu Look at the periodic table, and tell whether the second isotope is 63Cu or 66Cu 2.117 Sulfur has four naturally occurring isotopes, including 33S, RAN 34S, and 36S Look at the periodic table, and tell whether the fourth isotope is 32S or 35S 2.118 Naturally occurring boron consists of two isotopes: 10B (19.9%) with an isotopic mass of 10.0129 and 11B (80.1%) with an isotopic mass of 11.009 31 What is the atomic weight of boron? Check your answer by looking at a periodic table 2.119 Naturally occurring silver consists of two isotopes: 107Ag (51.84%) with an isotopic mass of 106.9051 and 109Ag (48.16%) with an isotopic mass of 108.9048 What is the atomic weight of silver? Check your answer in a periodic table 2.120 Magnesium has three naturally occurring isotopes: 24Mg (23.985) with 78.99% abundance, 25Mg (24.986) with 10.00% abundance, and a third with 11.01% abundance Look up the atomic weight of magnesium, and then calculate the mass of the third isotope 2.121 A sample of naturally occurring silicon consists of 28Si (27.9769), 29Si (28.9765), and 30Si (29.9738) If the atomic weight of silicon is 28.0855 and the natural abundance of 29Si is 4.68%, what are the natural abundances of 28Si and 30Si? 2.122 Copper metal has two naturally occurring isotopes: copper-63 RAN (69.15%; isotopic mass = 62.93) and copper-65 (30.85%; isotopic mass 64.93) Calculate the atomic weight of copper, and check your answer in the periodic table 2.123 Germanium has five naturally occurring isotopes: 70Ge, RAN 20.5%, 69.924; 72Ge, 27.4%, 71.922; 73Ge, 7.8%, 72.923; 74 Ge, 36.5%, 73.921; and 76Ge, 7.8%, 75.921 What is the atomic weight of germanium? 2.124 What is the mass in grams of each of the following samples? RAN (a) 1.505 mol of Ti (b) 0.337 mol of Na (c) 2.583 mol of U 80 chapter Atoms, Molecules, and Ions 100 90 80 Intensity 70 60 50 40 30 120 Intensity 60 (c) 90 80 70 60 50 40 30 20 10 63 65 Mass/charge 45.9527 11.19 30 70 100 Ion intensity 40 60 Mass/charge Ion mass/charge ratio 50 70 20 43 44 45 46 47 48 46.9518 10.09 47.9479 100.00 48.9479 7.43 49.9448 7.03 49 50 51 52 53 Mass/charge 50 Chemical Compounds (Sections 2.11–2.12) 40 124 60 65 123 70 80 63 122 80 80 20 10 121 90 10 20 10 74.80 100 90 30 100.00 122.9042 2.133 Use the data from the mass spectrum of a sample of an element to calculate the element’s atomic weight Identify the element 100 30 120.9038 Mass/charge 90 40 Ion intensity 10 100 50 Ion mass/charge ratio 20 (b) Intensity 2.132 Use the data from the mass spectrum of a sample of an element to calculate the element’s atomic weight Identify the element (a) Intensity Intensity 2.125 How many moles are in each of the following samples? (a) 11.51 g of Ti (b) 29.127 g of Na (c) 1.477 kg of U 2.126 If the atomic weight of an element is x, what is the mass in grams RAN of 6.02 * 1023 atoms of the element? How does your answer compare numerically with the atomic weight of element x? 2.127 If the atomic weight of an element is x, what is the mass in RAN grams of 3.17 * 1020 atoms of the element? 2.128 If 6.02 * 1023 atoms of element Y have a mass of 83.80 g, what is the identity of Y? 2.129 If 4.61 * 1021 atoms of element Z have a mass of 0.815 g, what is the identity of Z? 2.130 Refer to Figure 2.10 showing a schematic illustration of a mass spectrometer (a) What is the purpose of bombarding the gaseous atoms with an electron beam? (b) Compare two ions with a + charge traveling through the curved, evacuated tube in the mass spectrometer Will a heavier ion or lighter ion be deflected to a greater degree by the magnetic field? (c) Under a given set of experimental conditions ions of a certain mass-to-charge ratio pass through a slit and strike the detector What experimental variable in the mass spectrometer is altered so that a different mass-tocharge ratio ion strikes the detector? 2.131 Copper has two naturally occurring isotopes, 63Cu (relative abundance = 69.17%) and 65Cu (relative abundance = 30.83%) Select which mass spectrum represents a sample of copper 63 65 Mass/charge 2.134 What is the difference between a covalent bond and an ionic bond? 2.135 Which of the following bonds are likely to be covalent and which ionic? Explain (a) B ¬ Br (b) Na ¬ Br (c) Br ¬ Cl (d) O ¬ Br 2.136 The symbol CO stands for carbon monoxide, but the symRAN bol Co stands for the element cobalt Explain 2.137 Correct the error in each of the following statements: (a) The formula of ammonia is NH3 (b) Molecules of potassium chloride have the formula KCl (c) Cl- is a cation (d) CH4 is a polyatomic ion 2.138 How many protons and electrons are in each of the following ions? (a) Be2+ (b) Rb+ 2(c) Se (d) Au3+ 2.139 What is the identity of the element X in the following ions? (a) X2+, a cation that has 36 electrons (b) X-, an anion that has 36 electrons Section Problems 81 2.140 The structural formula of isopropyl alcohol, better known as “rubbing alcohol,” is shown What is the chemical formula of isopropyl alcohol? H H H O H C C C H H H H Isopropyl alcohol 2.141 Lactic acid, a compound found both in sour milk and in tired muscles, has the structure shown What is its chemical formula? H H H O C C C H O O H H Lactic acid 2.142 Butane, the fuel used in disposable lighters, has the formula C4H10 The carbon atoms are connected in the sequence C ¬ C ¬ C ¬ C, and each carbon has four covalent bonds Draw the structural formula of butane 2.143 Cyclohexane, C6H12, is an important starting material used in the industrial synthesis of nylon Each carbon has four covalent bonds, two to hydrogen and two to other carbons Draw the structural formula of cyclohexane 2.144 Isooctane, the substance in gasoline from which the term octane rating derives, has the formula C8H18 Each carbon has four covalent bonds, and the atoms are connected in the sequence shown Draw the complete structural formula of isooctane C C C C C C C C 2.145 Fructose, C6H12O6, is the sweetest naturally occurring sugar and is found in many fruits and berries Each carbon has four covalent bonds, each oxygen has two covalent bonds, each hydrogen has one covalent bond, and the atoms are connected in the sequence shown Draw the complete structural formula of fructose O C C C O O C O C C O O Naming Compounds (Section 2.13) 2.146 Give systematic names for the following binary compounds: (a) CsF (b) K2O (c) CuO 2.147 Give systematic names for the following binary compounds: (a) BaS (b) BeBr2 (c) FeCl3 2.148 Write formulas for the following binary compounds: (a) Potassium chloride (b) Tin(II) bromide (c) Calcium oxide (d) Barium chloride (e) Aluminum hydride 2.149 Write formulas for the following binary compounds: RAN (a) Vanadium(III) chloride (b) Manganese(IV) oxide (c) Copper(II) sulfide (d) Aluminum oxide 2.150 Write formulas for the following compounds: RAN (a) Calcium acetate (b) Iron(II) cyanide (c) Sodium dichromate (d) Chromium(III) sulfate (e) Mercury(II) perchlorate 2.151 Write formulas for the following compounds: RAN (a) Lithium phosphate (b) Magnesium hydrogen sulfate (c) Manganese(II) nitrate (d) Chromium(III) sulfate 2.152 Give systematic names for the following compounds: RAN (a) Ca(ClO)2 (b) Ag2S2O3 (c) NaH2PO4 (d) Sn(NO3)2 (e) Pb(CH3CO2)4 (f) (NH4)2SO4 2.153 Name the following ions: RAN (a) Ba2+ (b) Cs + 3+ (c) V (d) HCO3 (e) NH4+ (f) Ni2+ (g) NO2 (h) ClO22+ (i) Mn (j) ClO4 2.154 What are the formulas of the compounds formed from the RAN following ions? (a) Ca2+ and Br (b) Ca2+ and SO42(c) Al3+ and SO42 2.155 What are the formulas of the compounds formed from the RAN following ions? (a) Na+ and NO3(b) K+and SO422+ (c) Sr and Cl 2.156 Write formulas for compounds of calcium with each of the RAN following: (a) Chlorine (b) Oxygen (c) Sulfur 2.157 Write formulas for compounds of rubidium with each of the RAN following: (a) Bromine (b) Nitrogen (c) Selenium 2.158 Give the formulas and charges of the following ions: RAN (a) Sulfite ion (b) Phosphate ion (c) Zirconium(IV) ion (d) Chromate ion (e) Acetate ion (f) Thiosulfate ion 2.159 What are the charges on the positive ions in the following compounds? (a) Zn(CN)2 (b) Fe(NO2)3 (c) Ti(SO4)2 (d) Sn3(PO4)2 (e) (f) MnO2 Hg2S (g) KIO4 (h) Cu(CH3CO2)2 82 chapter Atoms, Molecules, and Ions 2.160 Name the following binary molecular compounds: RAN (a) CCl4 (b) ClO2 (c) N2O (d) N2O3 2.161 Give systematic names for the following compounds: RAN (a) NCl3 (b) P4O6 (c) S2F2 2.162 Name the following binary compounds of nitrogen and RAN oxygen: (a) NO (b) N2O (c) NO2 (d) N2O4 (e) N2O5 2.163 Name the following binary compounds of sulfur and oxygen: RAN (a) SO (b) S2O2 (c) S5O (d) S7O2 (e) SO3 2.164 Fill in the missing information to give formulas for the following compounds: (a) Na?SO4 (b) Ba?(PO4)? (c) Ga?(SO4)? 2.165 Write formulas for each of the following compounds: (a) Sodium peroxide (b) Aluminum bromide (c) Chromium(III) sulfate MULTICONCEPT PROBLEMS 2.166 Ammonia (NH3) and hydrazine (N2H4) are both compounds RAN of nitrogen and hydrogen Based on the law of multiple proportions, how many grams of hydrogen would you expect 2.34 g of nitrogen to combine with to yield ammonia? To yield hydrazine? 2.167 If 3.670 g of nitrogen combines with 0.5275 g of hydrogen RAN to yield compound X, how many grams of nitrogen would combine with 1.575 g of hydrogen to make the same compound? Is X ammonia (NH3) or hydrazine (N2H4)? 2.168 Prior to 1961, the atomic mass unit (amu) was defined as RAN 1/16 the mass of the atomic weight of oxygen; that is, the atomic weight of oxygen was defined as exactly 16 What was the mass of a 12C atom prior to 1961 if the atomic weight of oxygen on today’s scale is 15.9994? 2.169 What was the mass in atomic mass units of a 40Ca atom RAN prior to 1961 if its mass on today’s scale is 39.9626? (See Problem 2.168.) 2.170 The molecular weight of a compound is the sum of the RAN atomic masses of all atoms in the molecule What is the molecular mass of acetaminophen (C8H9NO2), the active ingredient in Tylenol? 2.171 The mass percent of an element in a compound is the mass RAN of the element (total mass of the element’s atoms in the compound) divided by the mass of the compound (total mass of all atoms in the compound) times 100% What is the mass percent of each element in acetaminophen? (See Problem 2.170.) Acetaminophen (C8H9NO2) 2.172 In an alternate universe, the smallest negatively charged RAN particle, analogous to our electron, is called a blorvek To determine the charge on a single blorvek, an experiment like Millikan’s with charged oil droplets was carried out, and the following results were recorded: Droplet Number Charge (C) 7.74 * 10 -16 4.42 * 10 -16 2.21 * 10 -16 4.98 * 10 -16 6.64 * 10 -16 (a) Based on these observations, what is the largest possible value for the charge on a blorvek? (b) Further experiments found a droplet with a charge of 5.81 * 10-16 C Does this new result change your answer to part (a)? If so, what is the new largest value for the blorvek’s charge? ... 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)... 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 Evers College Joe Casalnuovo,... 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