This International Student Edition is for use outside of the U.S Julia Burdge Michelle Driessen Introductory Chemistry AN ATOMS FIRST APPROACH Second Edition Fundamental Constants Avogadro’s number (NA) 6.0221418 × 1023 Electron charge (e) 1.6022 × 10−19 C Electron mass Faraday constant (F) Gas constant (R) 9.109387 × 10−28 g 96,485.3 C/mol e− 0.0821 L ⋅ atm/K ⋅ mol 8.314 J/K ⋅ mol 62.36 L ⋅ torr/K ⋅ mol 1.987 cal/K ⋅ mol Planck’s constant (h) 6.6256 × 10−34 J ⋅ s Proton mass 1.672623 × 10−24 g Neutron mass 1.674928 × 10−24 g Speed of light in a vacuum 2.99792458 × 108 m/s Some Prefixes Used with SI Units tera (T) 1012 centi (c) 10−2 giga (G) 109 milli (m) 10−3 mega (M) 106 micro ( µ) 10−6 kilo (k) 103 nano (n) 10−9 deci (d) 10−1 pico (p) 10−12 Useful Conversion Factors and Relationships lb = 453.6 g in = 2.54 cm (exactly) mi = 1.609 km km = 0.6215 mi pm = × 10−12 m = × 10−10 cm atm = 760 mmHg = 760 torr = 101,325 N/m2 = 101,325 Pa cal = 4.184 J (exactly) L ⋅ atm = 101.325 J 1J=1C×1V ?°C = (°F − 32°F) × ?°F = 5°C 9°F 9°F × (°C) + 32°F 5°C 1K ?K = (°C + 273.15°C) ( 1°C ) Na Mg K Rb Cs Fr Lanthanum 138.9 89 La Yttrium 88.91 57 Y Scandium 44.96 39 Radium (226) Metalloids Rf Cr Mn 25 7B Tc Actinides Ru Iron 55.85 44 Fe 26 Ta Db Tantalum 180.9 105 W Sg Tungsten 183.8 106 Re Bh Rhenium 186.2 107 58 Thorium 232.0 Th Cerium 140.1 90 Ce 59 61 Mt Pa Protactinium 231.0 U Uranium 238.0 62 Rg Gold 197.0 111 Au Silver 107.9 79 Ag Copper 63.55 47 29 Cu 64 Gd Cn Mercury 200.6 112 Hg Cadmium 112.4 80 Cd Zinc 65.41 48 30 Zn 2B 12 Terbium 158.9 97 65 Tb Curium (247) Ge Silicon 28.09 32 Si Carbon 12.01 14 As Phosphorus 30.97 33 P Nitrogen 14.01 15 Nh Thallium 204.4 113 Tl Indium 114.8 81 In Fl Lead 207.2 114 Pb Tin 118.7 82 Sn Mc Bismuth 209.0 115 Bi Antimony 121.8 83 Sb Gallium Germanium Arsenic 69.72 72.64 74.92 49 50 51 Ga Aluminum 26.98 31 Al Boron 10.81 13 N 5A 15 Lv Polonium (209) 116 Po Tellurium 127.6 84 Te Selenium 78.96 52 Se Sulfur 32.07 34 S Oxygen 16.00 16 O 6A 16 Ts Astatine (210) 117 At Iodine 126.9 85 I Bromine 79.90 53 Br Chlorine 35.45 35 Cl Fluorine 19.00 17 F 7A 17 67 Ho Cf Es Dysprosium Holmium 162.5 164.9 98 99 66 Dy Thulium 168.9 101 69 Ytterbium 173.0 102 70 Tm Yb Fm Md No Erbium 167.3 100 68 Er Berkelium Californium Einsteinium Fermium Mendelevium Nobelium (247) (251) (252) (257) (258) (259) Pu Am Cm Bk Europium Gadolinium 152.0 157.3 95 96 63 Eu Neptunium Plutonium Americium (237) (244) (243) Np Ds Platinum 195.1 110 Pt Palladium 106.4 78 Pd Nickel 58.69 46 28 Ni 10 1B 11 C B 4A 14 3A 13 Main group Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine (293) (293) (280) (285) (286) (289) (289) (276) (281) Nd Pm Sm 60 Hs Hassium (270) Ir Iridium 192.2 109 Praseodymium Neodymium Promethium Samarium 140.9 144.2 (145) 150.4 91 92 93 94 Pr Os Osmium 190.2 108 Rhodium 102.9 77 Rh Cobalt 58.93 45 Co 27 8B Average atomic mass Symbol Niobium Molybdenum Technetium Ruthenium (98) 101.1 92.91 95.94 74 73 76 75 Nb Mo Vanadium Chromium Manganese 54.94 50.94 52.00 41 42 43 V 24 6B An element Rutherfordium Dubnium Seaborgium Bohrium (267) (272) (268) (271) Lanthanides Actinium (227) Hafnium 178.5 104 Hf Zirconium 91.22 72 Zr Titanium 47.87 40 Ti 23 22 21 Sc 5B 4B C Carbon 12.01 Transition metals Name Atomic number Key Periodic Table of the Elements 3B Ra Ac Barium 137.3 88 Ba Strontium 87.62 56 Sr Calcium 40.08 38 Ca Magnesium 24.31 20 Nonmetals Metals Francium (223) Cesium 132.9 87 Rubidium 85.47 55 Potassium 39.10 37 Sodium 22.99 19 Beryllium 9.012 12 Lithium 6.941 11 Be 2A Group number Hydrogen 1.008 H 1A Li Period number Main group Lawrencium (262) Lr Lutetium 175.0 103 71 Lu Oganesson (294) Og Radon (222) 118 Rn Xenon 131.3 86 Xe Krypton 83.80 54 Kr Argon 39.95 36 Ar Neon 20.18 18 Ne Helium 4.003 10 He 8A 18 7 List of the Elements with Their Symbols and Atomic Masses* Element Actinium Aluminum Americium Antimony Argon Arsenic Astatine Barium Berkelium Beryllium Bismuth Bohrium Boron Bromine Cadmium Calcium Californium Carbon Cerium Cesium Chlorine Chromium Cobalt Copernicium Copper Curium Darmstadtium Dubnium Dysprosium Einsteinium Erbium Europium Fermium Flerovium Fluorine Francium Gadolinium Gallium Germanium Gold Hafnium Hassium Helium Holmium Hydrogen Indium Iodine Iridium Iron Krypton Lanthanum Lawrencium Lead Lithium Livermorium Lutetium Magnesium Manganese Meitnerium Symbol Atomic Number Atomic Mass† Ac Al Am Sb Ar As At Ba Bk Be Bi Bh B Br Cd Ca Cf C Ce Cs Cl Cr Co Cn Cu Cm Ds Db Dy Es Er Eu Fm Fl F Fr Gd Ga Ge Au Hf Hs He Ho H In I Ir Fe Kr La Lr Pb Li Lv Lu Mg Mn Mt 89 13 95 51 18 33 85 56 97 83 107 35 48 20 98 58 55 17 24 27 112 29 96 110 105 66 99 68 63 100 114 87 64 31 32 79 72 108 67 49 53 77 26 36 57 103 82 116 71 12 25 109 (227) 26.9815386 (243) 121.760 39.948 74.92160 (210) 137.327 (247) 9.012182 208.98040 (272) 10.811 79.904 112.411 40.078 (251) 12.0107 140.116 132.9054519 35.453 51.9961 58.933195 (285) 63.546 (247) (281) (268) 162.500 (252) 167.259 151.964 (257) (289) 18.9984032 (223) 157.25 69.723 72.64 196.966569 178.49 (270) 4.002602 164.93032 1.00794 114.818 126.90447 192.217 55.845 83.798 138.90547 (262) 207.2 6.941 (293) 174.967 24.3050 54.938045 (276) Element Mendelevium Mercury Molybdenum Moscovium Neodymium Neon Neptunium Nickel Nihonium Niobium Nitrogen Nobelium Oganesson Osmium Oxygen Palladium Phosphorus Platinum Plutonium Polonium Potassium Praseodymium Promethium Protactinium Radium Radon Rhenium Rhodium Roentgenium Rubidium Ruthenium Rutherfordium Samarium Scandium Seaborgium Selenium Silicon Silver Sodium Strontium Sulfur Tantalum Technetium Tellurium Tennessine Terbium Thallium Thorium Thulium Tin Titanium Tungsten Uranium Vanadium Xenon Ytterbium Yttrium Zinc Zirconium Symbol Atomic Number Md Hg Mo Mc Nd Ne Np Ni Nh Nb N No Og Os O Pd P Pt Pu Po K Pr Pm Pa Ra Rn Re Rh Rg Rb Ru Rf Sm Sc Sg Se Si Ag Na Sr S Ta Tc Te Ts Tb Tl Th Tm Sn Ti W U V Xe Yb Y Zn Zr 101 80 42 115 60 10 93 28 113 41 102 118 76 46 15 78 94 84 19 59 61 91 88 86 75 45 111 37 44 104 62 21 106 34 14 47 11 38 16 73 43 52 117 65 81 90 69 50 22 74 92 23 54 70 39 30 40 Atomic Mass† (258) 200.59 95.94 (289) 144.242 20.1797 (237) 58.6934 (286) 92.90638 14.0067 (259) (294) 190.23 15.9994 106.42 30.973762 195.084 (244) (209) 39.0983 140.90765 (145) 231.03588 (226) (222) 186.207 102.90550 (280) 85.4678 101.07 (267) 150.36 44.955912 (271) 78.96 28.0855 107.8682 22.98976928 87.62 32.065 180.94788 (98) 127.60 (293) 158.92535 204.3833 232.03806 168.93421 118.710 47.867 183.84 238.02891 50.9415 131.293 173.04 88.90585 65.409 91.224 *These atomic masses show as many significant figures as are known for each element The atomic masses in the periodic table are shown to four significant figures, which is sufficient for solving the problems in this book †Approximate values of atomic masses for radioactive elements are given in parentheses Introductory Chemistry An Atoms First Approach SECOND EDITION Julia Burdge COLLEGE OF WESTERN IDAHO Michelle Driessen UNIVERSITY OF MINNESOTA INTRODUCTORY CHEMISTRY Published by McGraw-Hill Education, Penn Plaza, New York, NY 10121 Copyright © 2020 by McGraw-Hill Education All rights reserved Printed in the United States of America No part of this publication may be reproduced or distributed in any form or by any means, or stored in a database or retrieval system, without the prior written consent of McGraw-Hill Education, including, but not limited to, in any network or other electronic storage or transmission, or broadcast for distance learning Some ancillaries, including electronic and print components, may not be available to customers outside the United States This book is printed on acid-free paper LWI 21 20 19 ISBN 978-1-260-56586-7 MHID 1-260-56586-6 Cover Image: ©ketkarn sakultap/Getty Images All credits appearing on page or at the end of the book are considered to be an extension of the copyright page Design Icon Credits: Animation icon: ©McGraw-Hill Education; Hot Spot Icon: ©LovArt/Shutterstock.com The Internet addresses listed in the text were accurate at the time of publication The inclusion of a website does not indicate an endorsement by the authors or McGraw-Hill Education, and McGraw-Hill Education does not guarantee the accuracy of the information presented at these sites mheducation.com/highered To the people who will always matter the most: Katie, Beau, and Sam —Julia Burdge To my family, the center of my universe and happiness, with special thanks to my husband for his support and making me the person I am today —Michelle Driessen And in memory of Raymond Chang He was a brilliant educator, a prolific writer, an extraordinary mentor, and a dear friend —Julia Burdge and Michelle Driessen About the Authors Julia Burdge holds a Ph.D (1994) from The University of Idaho in Moscow, Idaho; and a Master’s Degree from The University of South Florida Her research interests have included synthesis and characterization of cisplatin analogues, and development of new analytical techniques and instrumentation for measuring ultra-trace levels of atmospheric sulfur compounds ©David Spurgeon She currently holds an adjunct faculty position at The College of Western Idaho in Nampa, Idaho, where she teaches general chemistry using an atoms first approach; but spent the lion’s share of her academic career at The University of Akron in Akron, Ohio, as director of the Introductory Chemistry program In addition to directing the general chemistry program and supervising the teaching activities of graduate students, Julia established a future-faculty development program and served as a mentor for graduate students and postdoctoral associates Julia relocated back to the Northwest to be near family In her free time, she enjoys precious time with her three children, and with Erik Nelson, her husband and best friend Michelle Driessen earned a Ph.D in 1997 from the University of Iowa in Iowa City, Iowa Her research and dissertation focused on the thermal and photochemical reactions of small molecules at the surfaces of metal nanoparticles and high surface area oxides Following graduation, she held a tenure-track teaching and research position Courtesy of Michelle Driessen at Southwest Missouri State University for several years A family move took her back to her home state of Minnesota where she held positions as adjunct faculty at both St Cloud State University and the University of Minnesota It was during these adjunct appointments that she became very interested in chemical education Over the past several years she has transitioned the general chemistry laboratories at the University of Minnesota from verification to problem-based, and has developed both online and hybrid sections of general chemistry lecture courses She is currently the Director of General Chemistry at the University of Minnesota where she runs the general chemistry laboratories, trains and supervises teaching assistants, and continues to experiment with active learning methods in her classroom Michelle and her husband love the outdoors and their rural roots They take every opportunity to visit their family, farm, and horses in rural Minnesota viii Brief Contents 10 11 12 13 14 15 16 17 Atoms and Elements  Electrons and the Periodic Table  30 Compounds and Chemical Bonds  74 How Chemists Use Numbers  122 The Mole and Chemical Formulas  164 Molecular Shape  196 Solids, Liquids, and Phase Changes  238 Gases  272 Physical Properties of Solutions  312 Chemical Reactions and Chemical Equations  348 Using Balanced Chemical Equations  386 Acids and Bases  420 Equilibrium  458 Organic Chemistry  484 Biochemistry  510 Nuclear Chemistry  526 Electrochemistry  542 Appendix  Mathematical Operations  A-1 Glossary  G-1 Answers to Odd-Numbered Problems  AP-1 Index  I-1 SECTION 2.7   Ions: The Loss and Gain of Electrons 61 CHECKPOINT–SECTION 2.6  Periodic Trends 2.6.1 Arrange the elements Ca, Sr, and Ba in order of increasing amount of energy required to remove an electron a) Ca < Sr < Ba d) Sr < Ba < Ca b) Ba < Sr < Ca e) Sr < Ca < Ba a) K, Br b) Br, K c) Ba < Ca < Sr c) K, K 2.6.2 For each of the following pairs of elements, indicate which will have the greater tendency to gain an electron: Rb or Sr, C or N, O or F a) Rb, C, O d) Sr, N, O b) Sr, N, F e) Rb, C, F 2.6.3 Which element, K or Br, will have the greater tendency to lose an electron and which will have the greater tendency to gain an electron? d) Br, Br e) Both will have the same tendencies c) Sr, C, F 2.7   Ions: The Loss and Gain of Electrons In the last section, we discussed trends in the ability to lose an electron or to gain an electron In fact some atoms can lose more than one electron, and some can gain more When an atom loses or gains one or more electrons, it becomes an ion, specifically, an atomic ion, which is simply an atom in which the number of electrons is not equal to the number of protons Because an electron has a negative charge, the loss or gain of electrons causes an atom to become charged An atom that loses one or more electrons becomes positively charged, and we refer to it as a cation (pronounced cat-ion) An atom that gains one or more electrons becomes negatively charged, and we refer to it as an anion (pronounced an-ion) Let’s look at some specific examples of ions and how we represent them Electron Configuration of Ions Sodium, with atomic number 11, is at the far left of the periodic table in Group (1A) According to the trend in how easily atoms lose electrons, we expect sodium to lose an electron easily—and it does In fact, sodium loses an electron so readily that it is never found in nature in its elemental (neutral) form Rather, all of the sodium atoms found in nature are sodium ions, each having lost an electron The sodium ion is represented as Na+, with the charge indicated by a superscript Recall that the orbital diagram of sodium is: 1s 2s 2p 3s And we can write its electron configuration as: [Ne]3s1 When a sodium atom loses an electron to become a sodium ion, the electron lost is sodium’s only valence electron—in the 3s orbital With the loss of the 3s electron, the orbital diagram becomes: 1s 2s 2p And the electron configuration becomes simply: [Ne] 62 CHAPTER 2  Electrons and the Periodic Table Chlorine, with atomic number 17, is at the far right of the periodic table in Group 17 (7A) According to Figure 2.24, we expect chlorine to gain an electron readily—and it does The orbital diagram of the chlorine atom is: 1s 2s 2p 3s 3p And its electron configuration can be written as: 1s22s22p63s23p5 or [Ne]3s23p5 When an atom gains an electron, the electron goes into the lowest energy orbital available In the case of Cl, that’s the 3p orbital that has just one electron in it When a Cl atom has gained an electron, it becomes the ion Cl−, which, like the sodium ion, is written with a superscript charge The orbital diagram of the Cl− ion is: 1s 2s 2p 3s 3p And its electron configuration is [Ne]3s23p6 Li+ Be2+ Na+ Mg2+ K+ Ca2+ Rb+ Sr2+ N3– Al3+ P3– which is the same as the electron configuration for argon: [Ar] Now let’s take a moment to contemplate the significance of the electron configurations of these ions In each case, the resulting electron configuration is identical to the electron configuration of a noble gas The way chemists say this is that each ion is isoelectronic with a noble gas Na+ is isoelectronic with Ne, and Cl− is isoelectronic with Ar Recall that the trend in ease of electron loss (electrons being harder to remove as we go from left to right across a period) indicates that the element with the smallest tendency to gain an electron in each period is the noble gas Recall also that the noble gases [Group 18 (8A)] were not included in our discussion of how easily atoms gain electrons This is because the noble gases neither lose nor gain electrons readily The electron configurations of the noble gases are inherently stable—so there is no benefit to a noble gas atom from either losing or gaining electrons But for most main-group elements, achieving the electron configuration of a noble gas significantly increases their stability Both sodium and chlorine are very unstable as neutral atoms But when each has achieved a noble gas electron configuration (sodium by losing an electron and chlorine by gaining one), they become the highly stable ions, Na+ and Cl−, that make up the familiar, stable substance most of us sprinkle on our food every day: salt Knowing that atoms of the main-group elements will lose or gain electrons to achieve a noble gas electron configuration—to become isoelectronic with a noble gas—enables us to predict the number of electrons that will be lost or gained in each case—and therefore to O2– F– predict the charge on the resulting ions Calcium, for example, atomic number 20, will lose its two valence electrons to become isoelectronic S2– Cl– with argon: Se2– Br– Ca   [Ar]4s2 I– Ca2+   [Ar] Cs+ Ba2+ Oxygen will gain two electrons to become isoelectronic with neon: O [He]2s22p4 O2−   [He]2s22p6 or [Ne] Figure 2.25  Common ions of main-group elements Figure 2.25 shows the predictable charges on common ions of the maingroup elements SECTION 2.7   Ions: The Loss and Gain of Electrons 63 Sample Problem 2.10 lets you practice writing electron configurations for ions of main-group elements SAMPLE PROBLEM 2.10 Writing Electron Configurations for Common Ions of Main-Group Elements Predict the charge on the common ion formed from each element listed and write the electron configuration for the ion (a) Na   (b) Ca   (c) O Strategy  Locate the element on the periodic table and note that main-group elements gain or lose electrons to take on the electron configuration of the nearest noble gas Setup  (a) Sodium has 11 protons and 11 electrons as a neutral atom Sodium will take on the electron configuration of Ne when it becomes an ion by losing one electron (b) Calcium has 20 protons and 20 electrons as a neutral atom It will take on the electron configuration of argon when it becomes an ion by losing electrons (c) Oxygen has protons and electrons as a neutral atom It will take on the electron configuration of neon when it becomes an  ion by gaining electrons Solution  (a) The sodium ion contains 11 protons and 10 electrons, leaving one “extra” positive charge for a +1 charge Its electron configuration is the same as that of neon, [He]2s22p6 (b) The calcium ion contains 20 protons and 18 electrons (after the loss of two), leaving an “extra” two positive charges for a +2  charge The electron configuration for the calcium ion is the same as that of argon, [Ne]3s23p6 (c) The oxygen ion contains protons and 10 electrons (after the gain of two), leaving an “extra” two negative charges for a –2  charge The electron configuration for the oxygen ion is the same as that of neon, [He]2s22p6 TH IN K A BO U T IT Nonmetals gain electrons to take on the electron configuration of the noble gas to their right Metals lose electrons to take on the electron configuration of the noble gas in the row above them Practice Problem A TTEMPT   Predict the charge on the common ion formed from each element listed and write the electron configuration for the ion (a) Br   (b) K   (c) S Practice Problem B UILD   Given the atomic number (Z) and the electron configuration, identify the ion represented (a) Z = 34, [Ar]4s23d104p6   (b) Z = 35, [Ar]4s23d104p6   (c) Z = 19, [Ne]3s23p6 Practice Problem C ONCEPTUALIZE   Would it be possible to identify the ions listed in Practice Problem B if you were not given the atomic number? Why or why not? Lewis Dot Symbols of Ions Just as we can represent atoms of main-group elements with Lewis dot symbols, showing each valence electron as a dot, we can represent ions of main-group elements with Lewis dot symbols To this, we start with the Lewis dot symbol of an atom and remove a dot for each electron it loses—or add a dot for each electron it gains Using the examples of sodium, which loses one electron to become the cation Na+, and chlorine, which gains one electron to become the anion Cl−, the Lewis dot symbols are Na+ and Cl – Note that there are no dots around the symbol for sodium As a neutral atom, sodium has just one valence electron It loses its only valence electron to become the ion Na+, which has the same electron configuration as the noble gas neon (Ne) There are eight dots around the symbol for chlorine As a neutral atom, chlorine has seven valence electrons To become the Cl− ion, it gains one electron and achieves the electron configuration of the noble gas argon (Ar) Note also that the Lewis symbol for the anion is enclosed in square brackets, with the charge shown outside the brackets 64 CHAPTER 2  Electrons and the Periodic Table Sample Problem 2.11 lets you practice writing Lewis dot symbols for common main-group ions SAMPLE PROBLEM 2.11 Writing Lewis Dot Symbols for Common Ions of Main-Group Elements Write the Lewis dot symbols for the atom and the ion commonly formed by each element (a) P   (b) Se   (c) Sr Strategy  Locate the element on the periodic table and determine how many valence electrons it has Each valence electron is represented by a dot around the element symbol To determine the Lewis dot symbol of the ion, decide if the atom will gain or lose electrons to become like its nearest noble gas 3– (a) Pto become and P Setup  (a) P has valence electrons as an atom, and will gain more an ion with the electron configuration of argon (b) Se has valence electrons as an atom and will gain 3– more to become an ion with2– the electron configuration of krypton (a) P and P (b) Se and Se (c) Sr and Sr (b) Se and Se (c) Sr and Sr (c) Sr has valence electrons as an atom and will lose both to become an ion with the electron configuration of krypton Solution  (a) P (b) Se and P and Se 3– T H I N K ABOU T IT 2– 2– 2+ 2+ 2+ You should that negatively charged monatomic ions have the same Lewis dot symbol containing valence (c) notice Sr and Sr electrons You can see a similar pattern for positively charged monatomic ions in Practice Problem A Practice Problem A TTEMPT Write the Lewis dot symbols for the atom and the ion commonly formed by each element (a) K   (b) Mg   (c) Al Practice Problem B UILD X n– Given this generic Lewis dot symbol, the value of n, and the period where the element is located, determine the identity of the ion described (a) n = 2, period 4   (b) n = 3, period 2   (c) n = 2, period Practice Problem C ONCEPTUALIZE Xn+ Given this generic Lewis dot symbol and the group the element belongs to, identify the charge (value of n) (a) Group (1A)   (b) Group (2A) We have seen that atoms have periodic variation in their ability to lose or gain electrons In fact, some atoms have no appreciable tendency to either In Chapter 3, we will examine the results of these tendencies and how they govern the interactions of atoms and give rise to most of the matter we encounter every day CHECKPOINT–SECTION 2.7  Ions: The Loss and Gain of Electrons 2.7.1 Which of the following ions are isoelectronic with a noble gas? (Select all that apply.) a) K2+ c) Br− 2+ 2+ b) Ca e) F− d) O c) I− and Kr b) O2− and Mg2+ d) S2− and Cl− a) [Ne]3p4 d) [Ne]3p6 2 b) [Ne]3s 3p 2.7.2 Which of the following pairs are isoelectronic with each other? (Select all that apply.) a) Ca2+ and Sr2+ 2.7.3 Select the correct ground-state electron configuration for S2− e) He and H+ c) [Ne]3s 3p e) [Ne] Key Terms 65 Chapter Summary Section 2.1 Section 2.4 ∙ Light has been very important in the experiments that led to the development of our current model of the atom Light is a form of energy that has certain characteristics, including wavelength (λ) and frequency (ν) ∙ An electron configuration specifies the arrangement of an atom’s electrons in its atomic orbitals Each atomic orbital can accommodate a maximum of two electrons Two electrons occupying the same atomic orbital must have opposite spin, meaning that they are spinning in opposite directions This is the Pauli exclusion principle Hund’s rule states that electrons in an atom in the ground state will occupy the lowest energy orbitals possible ∙ The familiar rainbow is actually the sun’s visible emission spectrum, meaning that those are the visible wavelengths of light emitted by the sun What we usually refer to as “light” is actually the visible portion of the electromagnetic spectrum ∙ When atoms of an element such as hydrogen are excited by the addition of energy, they emit a characteristic set of discrete wavelengths known as a line spectrum Section 2.2 ∙ Experiments with light led scientists to propose that energy (specifically light) is quantized, meaning that it consists of tiny, discrete “packages,” rather than being continuous The Bohr atom model explained the line spectrum of hydrogen with the concepts of electrons orbiting the nucleus like planets around the sun In the Bohr model, an orbit was defined by a quantum number (n) The lowest energy orbit (n = 1) was the ground state, and any orbit of higher energy (n > 1) was called an excited state Each of the four lines in the hydrogen line spectrum was the result of an electron transition from a higher orbit (n = 3, 4, 5, or 6) to a lower orbit (n = 2) Section 2.3 ∙ The most current model of the atom is the quantum mechanical (QM) model The QM model does not specify the location of an electron within an atom, but rather allows calculation of the probability of finding an electron within a region of space called an atomic orbital To define an atomic orbital requires a principal quantum number (n) (designated with an integer: 1, 2, 3, 4…), which specifies the principal energy level, and a letter designation (s, p, d, or f ), which specifies the energy sublevel Section 2.5 ∙ The electron configuration of an atom can be determined using the periodic table Atoms can be represented with Lewis dot symbols, which consist of the element’s symbol surrounded by up to eight dots—each representing an electron in the atom’s outermost s and p energy sublevels Section 2.6 ∙ Elements exhibit periodic trends in certain properties including atomic size, metallic character, ionization energy (how easily an atom loses an electron), and electron affinity (how easily an atom gains an electron) Section 2.7 ∙ When an atom either loses or gains one or more electrons, it becomes an atomic ion An atom that loses one or more electrons becomes positively charged and is called a cation An atom that gains one or more electrons becomes negatively charged and is called an anion Atoms of main-group elements typically lose or gain enough electrons to become isoelectronic with a noble gas Isoelectronic species have the same number of electrons This enables us to predict the charges on the ions of main-group elements Key Terms Anion  61 Emission spectrum  32 Line spectra  33 Quantization  34 Atomic ion  61 Energy sublevel  41 Metallic character  57 Quantum number, n  34 Atomic orbital  41 Excited state  35 Orbit  34 Spin  46 Bohr atom  34 Frequency (ν)  31 Pauli exclusion principle  47 Valence electrons  53 Cation  61 Ground state  35 Photon  34 Wavelength (λ)  31 Core electrons  53 Hund’s rule  47 Principal energy level  41 Electromagnetic spectrum  32 Isoelectronic  62 Principal quantum number, n  41 Electron configuration  46 Lewis dot symbol  54 QM model  41 Determining Ground-State Valence Electron Configurations Using the Periodic Table KEY SKILLS An easy way to determine the electron configuration of an element is by using the periodic table Although the table is arranged by atomic number, it is also divided into blocks that indicate the type of orbital occupied by an element’s outermost electrons Outermost valence electrons of elements in the s block (shown below in yellow), reside in s orbitals; those of elements in the p block (blue) reside in p orbitals; and so on (a) 1s 2s 2p 1s 2s 2p 3s 1s 2s 2p 3s (b) (c) 3p To determine the ground-state electron configuration of any element, we start with the most recently completed noble gas core and count across the following period to determine the valence electron configuration Consider the example of Cl, which has atomic number 17 The noble gas that precedes Cl is Ne, with atomic number 10 Therefore, we begin by writing [Ne] The noble gas symbol in square brackets represents the core electrons with a completed p subshell To complete the electron configuration, we count from the left of period as shown by the red arrow, adding the last (rightmost) configuration label from each block the arrow touches: 1s 2s 11 3s 19 4s 21 5s 55 6s 87 7s 1s 12 13 20 21 22 39 56 89 4d2 72 57 3d2 40 88 22 5d2 104 6d2 58 90 23 24 41 25 42 73 59 60 91 92 4f 5f 62 93 94 [Ne] 66 63 3s 65 96 111 64 95 10 49 10 66 81 10 112 113 10 66 97 31 48 79 110 30 47 78 109 61 46 77 108 29 107 45 76 28 75 106 44 27 43 74 105 26 67 98 3p5 10 99 10 14 15 32 33 50 51 82 83 114 115 2p2 3p2 4p2 5p2 6p2 7p2 68 11 100 11 16 17 34 35 52 53 84 85 116 69 12 101 12 117 70 71 13 13 118 102 86 54 36 18 10 14 103 14 There are seven electrons in addition to the noble gas core Two of them reside in an s subshell, and five of them reside in a p subshell By simply counting across the third period, we can determine the specific subshells that contain the valence electrons and arrive at the correct ground-state electron configuration: [Ne]3s23p5 For Ga, with atomic number 31, the preceding noble gas is Ar, with atomic number 18 Counting across the fourth period (green arrow) gives the ground-state electron configuration: [Ar]4s23d104p1 We can also use the periodic table to determine the identity of an element, given its ground-state electron configuration For example, given the configuration [Ne]3s23p4, we focus on the last entry in the configuration: 3p4 This tells us that the element is in the third period (3), in the p block ( p), and that it has four electrons in its p subshell (superscript 4) This corresponds to atomic number 16, which is the element sulfur (S) Key Skills Problems 2.1 Determine the atomic number of tin (Sn) using the periodic table on the inside cover of your book What is the noble gas core for Sn? a) Ar b) Kr c) Xe d) Ne e) Rn 2.4 Determine the atomic number of iodine (I) using the periodic table on the inside cover of your book What is the electron configuration of the I atom? a) [Xe]5s24d105p5 b) [Kr]5s25p5 c) [Xe]5s25p5 d) [Kr]5s24d105p5 e) [Kr]5s25d105p5 2.2 Determine the atomic number of vanadium (V) using the periodic table on the inside cover of your book Which of the following electron configurations correctly represents the V atom? a) [Ar]3d5 b) [Ar]4s23d2 c) [Ar]4s23d d) [Ar]4s23d3 e) [Kr]4s23d3 2.5 What is the electron configuration of the common ion that forms from the K atom? a) [Ar]4s1 b) [Ar] c) [Ar]4s2 d) [Kr] e) [Kr]4s1 2.3 What element is represented by the electron configuration [Kr]5s24d105p1? a) Sn b) Ga c) In d) Tl e) Zr 2.6 What common ions of main-group elements are isoelectronic with Ne? a) S2−, Cl−, and K+ b) O2−, F−, and Na+ c) F−, Cl−, and Br− d) O2−, S2−, and Se2− e) Li+, Na+, and K+ 68 CHAPTER 2  Electrons and the Periodic Table Questions and Problems SECTION 2.1: THE NATURE OF LIGHT 2.1 How are the wavelength and frequency of light related to one another? 2.2 How is the energy of light related to its frequency? 2.3 How is the energy of light related to its wavelength? 2.4 Which of the following colors of light has the longest wavelength: yellow, green, or violet? 2.5 Which of the following colors of light has the longest wavelength: red, blue, or orange? 2.6 Which of the following colors of light has the largest frequency: yellow, green, or violet? 2.7 Which of the following colors of light has the largest frequency: red, blue, or orange? 2.8 Place the following wavelengths of light in order of increasing frequency: 350 nm, 450 nm, and 550 nm 2.9 Place the following wavelengths of light in order of decreasing frequency: 400 nm, 550 nm, and 700 nm 2.10 Place the following types of electromagnetic radiation in order of increasing energy: infrared, X rays, and radio waves 2.11 Place the following types of electromagnetic radiation in order of increasing frequency: microwave, gamma rays, and visible 2.12 Place the following colors of light in order of increasing energy: violet, yellow, and red 2.13 Place the following colors of light in order of decreasing energy: blue, green, and orange SECTION 2.2: THE BOHR ATOM 2.14 Describe what the term quantized means in your own words 2.15 Define the term photon in your own words 2.16 What is meant by the term ground state when referring to an atom? 2.17 What is meant by the term excited state when referring to an atom? 2.18 Sketch the Bohr model of the hydrogen atom and include the following: (a) Label n = through n = (b) Label the absorption of light from n = (ground state) to n = (excited state) (c) Label the emission of light from n = to n = 2.19 Indicate which of the following electron transitions would be expected to emit visible light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.20 Indicate which of the following electron transitions would be expected to emit visible light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.21 Indicate which of the following electron transitions would be expected to emit any light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.22 Indicate which of the following electron transitions would be expected to emit any light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.23 Indicate which of the following electron transitions would be expected to absorb any light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.24 Indicate which of the following electron transitions would be expected to absorb any light in the Bohr model of the atom: (a) n = to n = (b) n = to n = (c) n = to n = 2.25 The hydrogen atom emission spectrum contains four wavelengths or colors of visible light Match the wavelength/color to the transition that it comes from Transition n = to n = n = to n = n = to n = n = to n = Wavelength/Color 657 nm/red 486 nm/green 434 nm/blue 410 nm/violet 2.26 How many photons of light are emitted during each of the following processes in a hydrogen atom? (a) one electron undergoes a transition from n = to n = (b) one electron undergoes a transition from n = to n = (c) twelve electrons (in 12 separate hydrogen atoms) undergo transitions from n = to n = (d) fifty electrons (in 50 separate hydrogen atoms) undergo transitions from n = to n = 69 Questions and Problems Visualizing Chemistry Figure 2.7 Visualizing Chemistry–Bohr Atom VC 2.1 Which of the following best explains why we see only four lines in the emission spectrum of hydrogen? a) Hydrogen has only four different electronic transitions b) Only four of hydrogen’s electronic transitions correspond to visible wavelengths c) The other lines in hydrogen’s emission spectrum can’t be seen easily against the black background VC 2.2 One way to see the emission spectrum of hydrogen is to view the hydrogen in an electric discharge tube through a spectroscope, a device that separates the wavelengths Why can we not view the emission spectrum simply by pointing the spectroscope at a sample of hydrogen confined in a glass tube or flask? a) Without the electrons being in excited states, there would be no emission of light b) The glass would make it impossible to see the emission spectrum c) Hydrogen alone does not exhibit an emission spectrum—it must be combined with oxygen VC 2.3 How many lines would we see in the emission spectrum of hydrogen if the downward transitions from excited states all ended at n = and no transitions ended at n = 2? a) We would still see four lines b) We would see five lines c) We would not see any lines VC 2.4 For a hydrogen atom in which the electron has been excited to n = 4, how many different transitions can occur as the electron eventually returns to the ground state? a) b) c) 2.34 2.35 2.36 2.37 2.38 2.39 2.40 2.41 2.42 SECTION 2.3: ATOMIC ORBITALS 2.27 2.28 2.29 2.30 2.31 2.32 2.33 What does an atomic orbital represent? How is a sublevel different from an orbital? Sketch the 3s, 3p, and 3d orbitals How is the 2s orbital different from the 3s orbital you sketched in Problem 2.29? How is it similar? How is a 4p orbital different from the 3p orbital you sketched in Problem 2.29? How is it similar? How is a 4d orbital different from the 3d orbital you sketched in Problem 2.29? How is it similar? Choose the larger orbital of each given pair (a) 2s or 4s (b) 2px or 2py (c) 3p or 4p 2.43 Choose the larger orbital of each given pair (a) 1s or 5s (b) 3pz or 4py (c) 3p or 4p Choose the smaller orbital of each given pair (a) 3d or 4d (b) 1s or 2s (c) 2px or 3px Choose the smaller orbital of each given pair (a) 3s or 4s (b) 5p or 2p (c) 4dy or 3dy Consider an electron in each of the following orbitals On average, which electron would be closer to the nucleus in each pair? (a) 1s or 3s (b) 2p or 4p (c) 3d or 4d Consider an electron in each of the following orbitals On average, which electron would be closer to the nucleus in each pair? (a) 2s or 2p (b) 3p or 5f (c) 4s or 4p Consider each of the following sublevel designations Determine if each is a legitimate designation and if not, explain why (a) 5p (c) 2f (b) 4s (d) 1p Consider each of the following sublevel designations Determine if each is a legitimate designation and if not, explain why (a) 1p (c) 3f (b) 6s (d) 4p Determine whether each of the following designations represents a single orbital, a sublevel, or both (a) 4d (c) 3px (b) 2s (d) 2p Determine whether each of the following designations represents a single orbital, a sublevel, or both (a) 5p (c) 2pz (b) 4s (d) 4p Match each of the following descriptions with the correct orbital designation Description Spherical orbital in the fourth level Dumbbell-shaped orbital in the second   level Cloverleaf-shaped orbital in the third   level Spherical orbital in the first level Orbital Designation 1s 2p 4s 3d 70 CHAPTER 2  Electrons and the Periodic Table 2.44 Place the following sublevels in order of decreasing energy (a) 3s, 3p, 3d (b) 1s, 2s, 3s (c) 2s, 2p, 3s 2.45 Place the following sublevels in order of decreasing energy (a) 4s, 4p, 4d (b) 2p, 3p, 4p (c) 1s, 2p, 3d SECTION 2.4: ELECTRON CONFIGURATIONS 2.46 What is meant by the term ground state when referring to electron configurations? 2.47 How many electrons can one orbital of any type hold? 2.48 How many electrons can be held in an orbital with the following designation? (a) s (c) d (b) p (d) f 2.49 How many orbitals does each subshell contain? (a) 2s (c) 2p (b) 4d (d) 4f 2.50 How many electrons can be held in a subshell with the following designation? (a) s (c) d (b) p (d) f 2.51 What is the maximum number of electrons that can be held in the following subshells? (a) 2p (c) 4s (b) 3d (d) 3p 2.52 Determine the ground-state electron configuration of the following atoms: (a) chlorine (b) calcium (c) nitrogen 2.53 Determine the ground-state electron configuration of the following atoms: (a) potassium (b) arsenic (c) selenium 2.54 Determine the ground-state electron configuration of the following atoms: (a) oxygen (b) sulfur (c) sodium 2.55 Determine the ground-state electron configuration of the following atoms: (a) lithium (b) silicon (c) magnesium 2.56 Draw the orbital diagram for each of the elements in Problem 2.52 2.57 Draw the orbital diagram for each of the elements in Problem 2.53 2.58 Draw the orbital diagram for each of the elements in Problem 2.54 2.59 Draw the orbital diagram for each of the elements in Problem 2.55 2.60 Write the condensed (using noble gas cores) ground-state electron configuration for the following atoms: (a) bromine (b) tellurium (c) cesium 2.61 Write the condensed (using noble gas cores) ground-state electron configuration for the following transition metals: (a) Fe (b) Zn (c) Ni 2.62 Write the condensed (using noble gas cores) ground-state electron configuration for the following transition metals: (a) Zr (b) Co (c) Mn 2.63 Write the condensed (using noble gas cores) ground-state electron configuration for the following transition metals: (a) Cd (b) Pd (c) V SECTION 2.5: ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE 2.64 Describe the difference between core and valence electrons 2.65 Determine the number of core electrons and valence electrons in each of the following elements: (a) phosphorus (c) calcium (b) iodine (d) potassium 2.66 Determine the number of core electrons and valence electrons in each of the following elements: (a) chlorine (c) cesium (b) nitrogen (d) arsenic 2.67 Label the diagram with the appropriate term describing each shaded “block.” I II III 2.68 Determine the ground-state electron configuration of the following atoms How many valence electrons does each have? (a) strontium (b) bromine (c) xenon Questions and Problems 2.69 Determine the ground-state electron configuration of the following atoms How many valence electrons does each have? (a) antimony (b) barium (c) tin 2.70 Determine the ground-state electron configuration of and draw the valence orbital diagram for each of the following atoms: (a) aluminum (c) phosphorus (b) rubidium (d) indium 2.71 Determine the ground-state electron configuration of and draw the valence orbital diagram for each of the following atoms: (a) iodine (c) krypton (b) selenium (d) strontium 2.72 How many unpaired electrons does each of the atoms in Problem 2.70 contain? 2.73 How many unpaired electrons does each of the atoms in Problem 2.71 contain? 2.74 Write out the ground-state electron configuration for the following atoms How many valence electrons does each have? What prediction can you make about the ions of these three elements? Is this related to their location on the periodic table? (a) lithium (b) sodium (c) potassium 2.75 Write out the ground-state electron configuration for the following atoms How many valence electrons does each have? What prediction can you make about the ions of these three elements? Is this related to their location on the periodic table? (a) fluorine (b) chlorine (c) bromine 2.76 Determine the element represented by each of the following ground-state electron configurations (a) [Ar]4s23d8 (b) [Ne]3s23p3 (c) [Ar]4s23d104p4 2.77 Determine the element represented by each of the following ground-state electron configurations (a) [Kr]5s24d105p2 (b) [Xe]6s1 (c) [Ar]4s23d104p5 2.78 Determine the identity of each element represented by the following orbital diagrams: (a) 1s 2s 2s 2s 2p2p2p2p2p 1s1s 1s 1s1s 1s1s 1s2s 2s 2s2s 2s 2s 2p2p2p2p 1s 1s1s 2s 2s2s 2p2p2p (b) 1s1s1s 1s1s 1s 3s 1s 1s2s 1s2s 2s2s 2p2p 2p2p 3s 3s3s 3s3s 2s2s 2s2s 2s 2p2p 2p2p 2p 3s3s 3s 1s 1s1s 2s 2s2s 2p2p2p 3s 3s3s (c) 1s1s 3s 1s 1s1s1s 1s 1s2s 2s2s 2s2s 2s 2p2p 2p2p2p 2p 3s3s 3s 3s3s 3s 1s2s 2s2s 2p2p 3s3s 1s 1s1s 2s 2s2s 2p2p2p 3s 3s3s 71 2.79 Determine the error(s) in the following groundstate electron configurations and identify the element represented (a) 1s22s12p53s23p24s2 (b) 1s32s22p1 (c) 1s22s52p23s1 2.80 Determine the error(s) in the following groundstate electron configurations and identify the element represented (a) 1s22p63s23p84s2 (b) 1s22s22p63p24s23d5 (c) 1s22s22p43s23p44s23d10 2.81 Which of the following elements would you expect to have properties most similar to F? Why? O    Ar    Ne    C    Br 2.82 Which of the following elements would you expect to have properties most similar to S? Why? Br    Cl    P    Se    Na 2.83 Draw the Lewis dot symbol for each of the following atoms: (a) magnesium (c) fluorine (b) phosphorus (d) argon 2.84 Draw Lewis dot symbols for the following atoms: (a) carbon (c) neon (b) chlorine (d) potassium 2.85 Draw Lewis dot symbols for the following atoms: (a) nitrogen (c) calcium (b) bromine (d) lithium 2.86 Draw the Lewis dot symbol for each of the following atoms What they have in common? How is this related to their location on the periodic table? (a) oxygen (b) sulfur (c) selenium 2.87 Draw the Lewis dot symbol for each of the following atoms What they have in common? How is this related to their location on the periodic table? (a) nitrogen (b) phosphorus (c) arsenic SECTION 2.6: PERIODIC TRENDS 2.88 What does it mean for an atom to lose an electron easily? 2.89 What types of elements gain electrons most easily? 2.90 Sketch an outline of the periodic table and draw a single arrow showing the overall trend of increasing atomic size 2.91 Sketch an outline of the periodic table and draw a single arrow showing the overall trend of increasing metallic character 2.92 Sketch an outline of the periodic table and draw a single arrow showing the overall trend of increasing difficulty of electron removal from an atom 3p3p 3p 3p 3p3p3p3p 3p3p3p 72 CHAPTER 2  Electrons and the Periodic Table 2.93 Sketch an outline of the periodic table and draw a single arrow showing the overall trend of increasing ease of electron gain 2.94 Based on the periodic trend, arrange the following sets of elements in order of decreasing atomic size (a) P, S, Cl (b) O, S, Se (c) Ca, K, Kr 2.95 Based on the periodic trend, arrange the following sets of elements in order of increasing atomic size (a) Rb, S, Sr (b) Ca, Li, Mg (c) K, Ca, Br 2.96 Can you place N, S, and Br in order of increasing atomic size based only on the periodic trend? Why or why not? 2.97 Select the element in each group that will lose an electron most easily, based on the periodic trend (a) F, O, S (b) Si, C, Ne (c) Li, Na, K 2.98 Select the element in each group from which it will be most difficult to remove an electron, based on the periodic trend (a) Ca, Se, As (b) S, Se, Ge (c) Li, C, O 2.99 Select the least metallic element in each group, based on the periodic trend (a) Si, S, K (b) Br, F, Ba (c) Na, Mg, P 2.100 Select the most metallic element in each group, based on the periodic trend (a) Cs, Ba, Sr (b) Li, Be, Na (c) K, Br, As SECTION 2.7: IONS: THE LOSS AND GAIN OF ELECTRONS 2.101 What is an ion? How is it different from an atom? 2.102 Describe a cation in your own words 2.103 Describe an anion in your own words 2.104 Determine the charge on each of the following ions: (a) Bromine with 36 electrons (b) Aluminum with 10 electrons (c) Zinc with 28 electrons (d) Arsenic with 36 electrons 2.105 Predict the common ion, including the charge, formed by each of the following elements: (a) magnesium (d) oxygen (b) potassium (e) iodine (c) phosphorus 2.106 Predict the common ion, including the charge, formed by each of the following elements: (a) barium (d) aluminum (b) nitrogen (e) tellurium (c) sulfur 2.107 Write the ground-state electron configuration for the common ion formed by each of the following elements: (a) nitrogen (b) strontium (c) bromine 2.108 Write the ground-state electron configuration for the common ion formed by each of the following elements: (a) chlorine (b) sulfur (c) calcium 2.109 Write the ground-state electron configuration for the common ion formed by each of the following elements: (a) rubidium (b) aluminum (c) phosphorus 2.110 Write the ground-state electron configuration for the common ion formed by each of the following elements What the ions have in common? How are they different from one another? (a) O (b) F (c) Na 2.111 Write the ground-state electron configuration for the common ion formed by each of the following elements What they have in common? How are they different from one another? (a) S (b) K (c) Ca 2.112 Write the ground-state electron configuration for the common ion formed by each of the following elements: (a) lithium (b) selenium (c) iodine 2.113 Draw the Lewis dot symbol for the common ion formed by each of the following elements: (a) Ca (d) O (b) K (e) N (c) F 2.114 Draw the Lewis dot symbol for the common ion formed by each of the following elements: (a) As (d) Li (b) S (e) Cs (c) Br ADDITIONAL PROBLEMS 2.115 A common salt substitute contains potassium chloride Draw the Lewis dot symbol for the potassium ion and the ion formed by chlorine 2.116 Sodium metal is highly reactive, but the sodium ion is stable Explain Answers to In-Chapter Materials FAMILIAR CHEMISTRY FIREWORKS 2.117 Copper is responsible for the blue color that is observed in fireworks Write the ground-state electron configuration and valence orbital diagram of copper The transition (movement of the electron between energy levels) in copper that is responsible for the blue light emitted is 4p to 4s Represent this transition using the orbital diagram you’ve drawn 2.118 Barium is responsible for the green color that is observed in fireworks Write the ground-state electron configuration and draw the valence orbital diagram of barium The transition (movement of the electron between energy levels) in barium that is responsible for this observed color is 6p to 6s Represent this transition using the orbital diagram you’ve drawn 2.119 Write the electron configuration and orbital diagram of the lithium atom If the 2s electron is excited to the 2p sublevel, and then subsequently returns to the ground state, it emits light with a wavelength of 670 nm, or 6.70 × 10−7 m 73 Determine the energy gap between the 2s and 2p sublevels in lithium (The relationship between energy and wavelength is expressed by the equation E = hc/λ, where h and c are constants with values of 6.626 ì 1034 J ã s and 3.00 ì 108 m/s, respectively; and λ is the wavelength of light expressed in meters.) 2.120 Write the electron configuration for the excited barium atom using information from Problem 2.118 2.121 Write the electron configuration for the excited lithium atom using information from Problem 2.119 2.122 A certain electron transition in an atom results in the emission of visible light If the difference in energy between the initial and final sublevels is 4.32 × 10−19 J, what is the wavelength of the emitted light? In what region of the visible spectrum does this wavelength fall? See the relationship between energy and wavelength in Problem 2.119 Answers to In-Chapter Materials Answers to Practice Problems 2.1A a, e 2.1B a, c, e 2.2A (a) No, the third principal energy level has no f sublevel; (b) yes; (c) yes; (d) no, the second principal energy level has no d sublevel 2.2B (a) 2, 3, and 4; (b) 1, 2, 3, and 4; (c) 4; (d) and 2.3A (a) 1s22s22p2, [He]2s22p2; (b) 1s22s22p6, [He]2s22p6; 1s22s22p63s23p5, [Ne]3s23p5 2.3B (a) Mg, (b) P 2.4A (a) [Kr]5s1, (b) [Ar]4s23d104p4, (c) [Kr]5s24d105p5 2.4B (a) Ge, (b) Sb, (c) Ba 2.5A (a) 5s1, (b) 4s24p4, (c) 5s25p5 2.5B (a) 4s24p2, (b) 5s25p3, (c) 6s2 2.6A (a) As, (b) I 2.6B (a) K > Se > Cl, (b) Sb > I > Br 2.7A (a) Mg, (b) N 2.7B (a) C < Li < K, (b) O < S < Si 2.8A (a) Mg, (b) Te 2.8B (a) C > Li > K, (b) As > Ca > Ba 2.9A (a) O, (b) N, (c) I 2.9B (a) Ga, (b) Sn, (c) Zn 2.10A (a) Br−, [Ar]4s23d104p6; (b) K+, [Ar]; (c) S2−, [Ne]3s23p6 2.10B (a) Se2−, (b) Br−, (c) K+ 2.11A (a) K, K+; (b) Mg , Mg2+; (c) Al , Al3+ 2.11B (a) Se2−, (b) N3−, (c) S2− Answers to Checkpoints 2.2.1 b 2.2.2 b, c, e 2.3.1 a 2.3.2 b, e 2.4.1 b 2.4.2 d 2.4.3 b 2.5.1 b 2.5.2 c 2.5.3 b, d, e 2.5.4 a, b, d 2.5.5 d 2.5.6 c 2.6.1 b 2.6.2 b 2.6.3 a 2.7.1 b, c, e 2.7.2 b, d 2.7.3 b CHAPTER Compounds and Chemical Bonds 3.1 Matter: Classification and Properties • States of Matter • Mixtures • Properties of Matter 3.2 3.3 Ionic Bonding and Binary Ionic Compounds Naming Ions and Binary Ionic Compounds • Naming Atomic Cations • Naming Atomic Anions • Naming Binary Ionic Compounds 3.4 Covalent Bonding and Molecules • Covalent Bonding • Molecules • Molecular Formulas 3.5 Naming Binary Molecular Compounds 3.6 Covalent Bonding in Ionic Species: Polyatomic Ions 3.7 Acids 3.8 Substances in Review Much of the matter we encounter exists in the form of mixtures Seawater is a mixture consisting predominantly of two familiar substances: water and salt âShutterstock/EpicStockMedia ã Distinguishing Elements and Compounds • Determining Whether a Compound Is Ionic or Molecular • Naming Compounds In This Chapter, You Will Learn How some atoms lose or gain electrons to become ions; and how some atoms combine to form molecules You will also learn about chemical bonding—the forces that hold atoms together in compounds; and about nomenclature—how to associate a compound’s formula with its name Things To Review Before You Begin • The charges on common main-group ions [ Figure 2.25] Most of the substances we encounter every day are not elements Rather, they are compounds—substances that consist of more than one element Two examples of compounds are water (H2O), which is a combination of the elements hydrogen and oxygen; and sodium chloride (NaCl), commonly known as salt, which is a combination of the elements sodium and chlorine The relationship between electron configuration and periodic properties of the elements that we learned about in Chapter gives us a way to understand the existence and the formation of compounds In this chapter, we explain what compounds are and how to name and identify them To this, we must begin by discussing matter in general 3.1   Matter: Classification and Properties Chemists classify matter as either a substance or a mixture of substances A substance is a form of matter that has a specific, universally constant composition and distinct properties, such as color, smell, and taste Familiar examples include water, salt, iron, mercury, carbon dioxide, and oxygen These examples differ from one another in composition, and we can tell them apart because they have different properties For example, salt can be dissolved in water; iron cannot Mercury is a liquid with a silvery appearance; carbon dioxide is a colorless gas Carbon dioxide can be used to extinguish a flame; oxygen, which is also a colorless gas, actually feeds a flame States of Matter All substances can, in principle, exist as a solid, a liquid, and a gas These three physical states are depicted in Figure 3.1 In a solid, particles are held close together in an orderly fashion with very little freedom of motion As a result, a solid does not conform to the shape of its container Particles in a liquid are also close together, but are not held rigidly in position; they can move around within the liquid This freedom of motion allows a liquid to conform to the shape of a container In a gas, the particles are separated by distances that are very large compared to the size of the particles A sample of gas conforms not only to the shape but also to the volume of its container We can convert a substance from one physical state to another without changing the identity of the substance For example, if we start with an ice cube and heat it, it will melt to form liquid water If we continue to heat the resulting liquid, it will boil, and vaporize to form a gas (water vapor) If we were to cool that water vapor, it would condense into liquid water Further cooling will cause it to freeze back into ice Despite the changes the water has undergone, it has remained water The identity of the substance has not changed despite our having converted it several times from one state to another Figure 3.2 illustrates the three physical states of water ... 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