Preview chemical principles the quest for insight, 7th edition by peter atkins (author), loretta jones (author), leroy laverman (author) (2016)

Preview Chemical Principles The Quest for Insight, 7th Edition by Peter Atkins (author), Loretta Jones (author), Leroy Laverman (author) (2016) Preview Chemical Principles The Quest for Insight, 7th Edition by Peter Atkins (author), Loretta Jones (author), Leroy Laverman (author) (2016) Preview Chemical Principles The Quest for Insight, 7th Edition by Peter Atkins (author), Loretta Jones (author), Leroy Laverman (author) (2016) Preview Chemical Principles The Quest for Insight, 7th Edition by Peter Atkins (author), Loretta Jones (author), Leroy Laverman (author) (2016) Preview Chemical Principles The Quest for Insight, 7th Edition by Peter Atkins (author), Loretta Jones (author), Leroy Laverman (author) (2016)

CHEMICAL PRINCIPLES THE QUEST FOR INSIGHT Seventh Edition PETER ATKINS LORETTA JONES LEROY LAVERMAN Periodic Table of the Elements H Period 1 Group 6.94 2s1 sodium Period 19 39.10 4s1 85.47 5s1 cesium 55 132.91 6s1 Fr 37 rubidium Cs 11 potassium Rb Be beryllium 9.01 2s2 22.99 3s1 K lithium Na 1.0079 1s1 Li hydrogen 87 francium (223) 7s1 Mg 12 magnesium 24.31 3s2 Ca calcium 20 40.08 4s2 Sr 38 strontium 87.62 5s2 Ba barium 56 137.33 6s2 Ra radium 88 (226) 7s2 Sc 21 scandium Ti yttrium 39 47.87 3d24s2 Zr 57 lanthanum 91.22 4d25s2 Hf 89 actinium 72 hafnium 138.91 5d16s2 Ac 40 zirconium 88.91 4d15s2 La 22 titanium 44.96 3d14s2 Y 178.49 5d26s2 Rf 104 rutherfordium (227) 6d17s2 (265) 6d27s2 Lanthanoids (lanthanides) Actinoids (actinides) Molar masses (atomic weights) quoted to the number of significant figures given here can be regarded as typical of most naturally occurring samples Elements 113, 115, 117, and 118 have been identified but not yet (in 2016) formally named V 23 vanadium 50.94 3d34s2 Nb niobium 92.91 4d45s1 Ta 73 tantalum 180.95 5d36s2 Db 105 dubnium (268) 6d37s2 Ce cerium 58 140.12 4f15d16s2 Th 41 thorium 90 232.04 6d27s2 Cr 24 chromium 52.00 3d54s1 Mo 42 molybdenum 95.94 4d55s1 W 74 tungsten 183.84 5d46s2 Sg 106 seaborgium (271) 6d47s2 Pr 59 praseodymium 140.91 4f36s2 Pa 91 protactinium Mn 25 manganese 54.94 3d54s2 Tc 43 technetium (98) 4d55s2 Re 75 rhenium 186.21 5d56s2 Bh 107 bohrium (272) 6d57s2 Nd 60 neodymium 144.24 4f46s2 U 92 uranium Fe 26 iron 55.84 3d64s2 Ru 44 ruthenium 101.07 4d75s1 Os osmium 76 190.23 5d66s2 Hs 108 hassium (270) 6d67s2 Pm 61 promethium (145) 4f56s2 Np 93 neptunium 231.04 238.03 (237) 5f26d17s2 5f36d17s2 5f46d17s2 Co 27 cobalt 58.93 3d74s2 Rh 45 rhodium 102.90 4d85s1 Ir iridium 77 192.22 5d76s2 Mt 109 meitnerium (276) 6d77s2 Sm 62 samarium 150.36 4f66s2 Pu 94 plutonium (244) 5f67s2 18 He helium 13 14 B Metal boron Metalloid 10.81 2s22p1 Nonmetal Al 13 aluminum 10 Ni nickel 11 28 58.69 3d84s2 Pd 46 palladium 106.42 4d10 Pt 78 platinum 195.08 5d96s1 Ds 110 darmstadtium (281) 6d87s2 Eu 63 europium 151.96 4f76s2 Am 95 americium (243) 5f77s2 Cu copper 29 63.55 3d104s1 Ag 47 silver 107.87 4d105s1 Au 79 gold 196.97 5d106s1 Rg 111 roentgenium (280) 6d107s1 Gd 64 gadolinium 157.25 4f75d16s2 Cm curium 26.98 3s23p1 12 96 (247) 5f76d17s2 Zn 30 zinc 65.41 3d104s2 Cd 48 cadmium 112.41 4d105s2 Hg 80 mercury 200.59 5d106s2 Cn Ga gallium 31 69.72 4s24p1 In indium 49 114.82 5s25p1 Tl thallium 81 204.38 6s26p1 112 113 copernicium (285) 6d107s2 Tb terbium 65 158.93 4f96s2 Bk 97 berkelium (247) 5f97s2 15 C carbon 12.01 2s22p2 Si 14 silicon 28.09 3s23p2 Ge 32 germanium 72.64 4s24p2 Sn 50 tin 118.71 5s25p2 Pb 82 lead 207.2 6s26p2 Fl 16 N nitrogen 14.01 2s22p3 P 15 phosphorus 30.97 3s23p3 As arsenic 33 74.92 4s24p3 Sb 51 antimony 121.76 5s25p3 Bi bismuth 83 208.98 6s26p3 114 115 flerovium (289) 7s27p2 Dy 66 dysprosium 162.50 4f106s2 Cf 98 californium (251) 5f107s2 Ho 67 holmium 164.93 4f116s2 Es 99 einsteinium (252) 5f117s2 17 O oxygen F 16 sulfur 19.00 2s22p5 Cl 34 selenium 35.45 3s23p5 Br 52 tellurium 127.60 5s25p4 Po 84 polonium (209) 6s26p4 Lv 35 bromine 78.96 4s24p4 Te 17 chlorine 32.06 3s23p4 Se fluorine 16.00 2s22p4 S 4.00 1s2 79.90 4s24p5 I iodine 53 126.90 5s25p5 At 85 astatine (210) 6s26p5 116 Ne 20.18 2s22p6 Ar 18 argon 39.95 3s23p6 Kr krypton 36 83.80 4s24p6 Xe xenon 54 131.29 5s25p6 Rn radon 86 (222) 6s26p6 117 livermorium 10 neon 118 (293) 7s27p4 Er erbium 68 167.26 4f126s2 Fm 100 fermium (257) 5f127s2 Tm thulium 69 168.93 4f136s2 Yb 173.04 4f146s2 Md 101 No mendelevium (258) 5f137s2 70 ytterbium 102 nobelium (259) 5f147s2 Lu 71 lutetium 174.97 5d16s2 Lr 103 lawrencium (262) 6d17s2 FREQUENTLY USED TABLES AND FIGURES Page Atomic and molecular properties Atomic radii Ionic radii First ionization energies Electron affinity Electronegativity Average bond lengths Ground-state electron configurations The elements (physical properties) Fig 1F.4 Fig 1F.6 Fig 1F.8 Fig 1F.12 Fig 2D.2 Table 2D.3 Appendix 2C Appendix 2D 54 55 57 59 97 101 A18 A19 Table 4C.1 Table 4E.1 Table 4E.3 Table 5A.2 Appendix 2A 268 291 293 351 A9 Table 6C.1 Table 6C.2 Table 6E.1 Table 6I.1 461 462 483 524 Table 6M.1 Appendix 2B 557 A16 Thermodynamic properties Standard enthalpies of physical change Lattice enthalpies Mean bond enthalpies Vapor pressure of water Thermodynamic data Solutions Acidity constants at 25 °C Basicity constants at 25 °C Acidity constants of polyprotic acids at 25 °C Solubility products Electrochemistry Standard potentials this'page'left'intentionally'blank THE QUEST FOR INSIGHT PETER ATKINS Oxford University LORETTA JONES University of Northern Colorado LEROY LAVERMAN University of California, Santa Barbara New York SEVENTH EDITION CHEMICAL PRINCIPLES Publisher: Kate Ahr Parker Library of Congress Control Number: 2015951706 Acquisitions Editor: Alicia Brady ISBN-13: 978-1-4641-8395-9 Developmental Editor: Heidi Bamatter ISBN-10: 1-4641-8395-3 Marketing Manager: Maureen Rachford Marketing Assistant: Cate McCaffery Media Editor: Amy Thorne Media Producer: Jenny Chiu Photo Editor: Robin Fadool Photo Licensing Editor: Richard Fox Senior Project Editor: Elizabeth Geller © 2016, 2013, 2010, 2005 by P W Atkins, L L Jones, and L E Laverman All rights reserved Printed in the United States of America First printing Cover Designer: Blake Logan International Edition Cover Design: Dirk Kaufman Text Designer: Marsha Cohen Art Manager: Matthew McAdams Illustrations: Peter Atkins and Leroy Laverman Production Manager: Susan Wein Composition: Aptara Printing and Binding: RR Donnelley Cover Image: © Ted Kinsman/Alamy W H Freeman and Company One New York Plaza Suite 4500 New York, NY 10004-1562 www.whfreeman.com Focus ATOMS Focus MOLECULES Focus STATES OF MATTER INTERLUDE Ceramics and Glasses Focus THERMODYNAMICS INTERLUDE Free Energy and Life Focus EQUILIBRIUM INTERLUDE Homeostasis Focus REACTIONS F1 67 145 239 241 346 347 442 INTERLUDE Practical Cells 443 584 Focus KINETICS 587 Focus THE MAIN-GROUP ELEMENTS 643 Focus THE d-BLOCK ELEMENTS 705 Focus 10 NUCLEAR CHEMISTRY 747 Focus 11 ORGANIC CHEMISTRY 777 829 INTERLUDE Technology: Fuels CONTENTS IN BRIEF FUNDAMENTALS MAJOR TECHNIQUES (Online Only) http://macmillanhighered.com/chemicalprinciples7e iii this'page'left'intentionally'blank xv F F.1 F.2 F.3 FUNDAMENTALS / F1 Introduction and Orientation A Matter and Energy A.1 A.2 A.3 A.4 B B.1 B.2 B.3 B.4 C.1 C.2 C.3 G.1 G.2 G.3 G.4 H.1 H.2 D.5 I F22 J J.2 J.3 E FUNDAMENTALS E Exercises / F44 F78 Oxidation and Reduction / F78 Oxidation Numbers / F80 How to Assign Oxidation Numbers / F80 Oxidizing and Reducing Agents / F82 Balancing Simple Redox Equations / F84 TOOLBOX K.1 K.3 K.4 Exercises / F37 The Mole / F38 Molar Mass / F40 Exercises / F77 Redox Reactions K.1 K.2 TOOLBOX D.2 How to Name Simple Inorganic Molecular Compounds / F33 The Nomenclature of Some Common Organic Compounds / F35 Moles and Molar Masses E.1 E.2 K F72 Acids and Bases in Aqueous Solution / F73 Strong and Weak Acids and Bases / F74 Neutralization / F76 FUNDAMENTALS J How to Name Ionic Compounds / F31 Names of Inorganic Molecular Compounds / F32 Exercises / F71 Acids and Bases J.1 F29 F66 Electrolytes / F66 Precipitates / F67 Ionic and Net Ionic Equations / F68 Putting Precipitation to Work / F69 FUNDAMENTALS I Exercises / F28 Exercises / F64 Precipitation Reactions I.1 I.2 I.3 I.4 Names of Cations / F29 Names of Anions / F29 Names of Ionic Compounds / F31 FUNDAMENTALS D F60 Representing Chemical Reactions / F60 Balanced Chemical Equations / F62 FUNDAMENTALS H TOOLBOX D.1 D.4 Exercises / F58 H Chemical Equations Exercises / F21 D Nomenclature F51 Classifying Mixtures / F51 Separation Techniques / F53 Concentration / F54 Dilution / F56 FUNDAMENTALS G F15 What Are Compounds? / F22 Molecules and Molecular Compounds / F23 Ions and Ionic Compounds / F24 FUNDAMENTALS C D.1 D.2 D.3 G Mixtures and Solutions Exercises / F13 Compounds Exercises / F50 TOOLBOX G.1 How to Calculate the Volume of Stock Solution Required for a Given Dilution / F57 Atoms / F15 The Nuclear Model / F16 Isotopes / F18 The Organization of the Elements / F19 FUNDAMENTALS B C F1 F5 Elements and Atoms F46 Mass Percentage Composition / F46 Determining Empirical Formulas / F48 Determining Molecular Formulas / F49 FUNDAMENTALS F Symbols and Units / F5 Accuracy and Precision / F8 Force / F9 Energy / F10 FUNDAMENTALS A The Determination of Composition CONTENTS Preface FUNDAMENTALS K Exercises / F85 F38 L Reaction Stoichiometry L.1 L.2 Mole-to-Mole Predictions / F87 Mass-to-Mass Predictions / F88 F87 v Topic 1E 49 Exercises What have you learned in this Topic? You have learned that the structures of many-electron atoms are explained by the systematic occupation of orbitals by electrons, with the order determined by the effects of penetration and shielding in conjunction with the Pauli exclusion principle The building-up principle is reflected in, and in a sense accounts for, the general structure of the periodic table The skills you have mastered are the ability to: h h Describe the factors affecting the energy of an electron in a many-electron atom (Section 1E.1) Write the ground-state electron configuration for an element (Toolbox 1E.1 and Example 1E.1) Topic 1E Exercises 1E.1 Which of the following increase when an electron in a lithium atom undergoes a transition from the 1s-orbital to a 2p-orbital? (a) Energy of the electron (b) Value of n (c) Value of l (d) Radius of the atom Which answers would be different for a hydrogen atom and in what way would they be different? (b) N (a) C 1s 2s 2p 2s 2p 1s 2s 2p (d) O (c) Be 1s 1E.2 Which of the following increase when an electron in a lith- 1s 2s 2p ium atom undergoes a transition from the 2s-orbital to a 2p-orbital? (a) Energy of the electron (b) Value of n (c) Value of l (d) Radius of the atom Which answers would be different for a hydrogen atom and in what way would they be different? 1E.8 Each of the following valence-shell configurations is possible for a neutral atom of a certain element What is the element and which configuration represents the ground state? 1E.3 (a) Write an expression for the total coulombic potential (a) energy for a lithium atom (b) What does each individual term represent? 1E.4 (a) Write an expression for the total coulombic potential energy for a beryllium atom (b) If Z denotes the number of electrons present in an atom, write a general expression to represent the total number of terms that will be present in the total coulombic potential energy expression 1E.5 Which of the following statements are true for many-electron atoms? If false, explain why (a) The effective nuclear charge Zeffe is independent of the number of electrons present in an atom (b) Electrons in an s-orbital are more effective than those in other orbitals at shielding other electrons from the nuclear charge because an electron in an s-orbital can penetrate to the nucleus of the atom (c) Electrons having l are better at shielding than electrons having l (d) Zeff e for an electron in a p-orbital is lower than for an electron in an s-orbital in the same shell 1E.6 For the electrons on a carbon atom in the ground state, decide which of the following statements are true If false, explain why (a) Zeff e for an electron in a 1s-orbital is the same as Zeff e for an electron in a 2s-orbital (b) Zeff e for an electron in a 2s-orbital is the same as Zeff e for an electron in a 2p-orbital (c) An electron in the 2s-orbital has the same energy as an electron in the 2p-orbital (d) The electrons in the 2p-orbitals have spin quantum numbers ms of opposite sign (e) The electrons in the 2s-orbital have the same value of the quantum number ms 1E.7 Determine whether each of the following electron configurations represents the ground state or an excited state of the atom given (b) 4s 4p 4s 4p 4s 4p 4s 4p (d) (c) 1E.9 Of the following sets of four quantum numbers {n, l, ml, ms}, identify the ones that are forbidden for an electron in an atom and explain why they are invalid: 1 (a) 54, 2, 21, 6; (b) 55, 0, 21, 12 6; (c) 54, 4, 21, 12 1E.10 Of the following sets of four quantum numbers {n, l, ml, ms}, identify the ones that are forbidden for an electron in an atom and explain why they are invalid: 1 (a) 52, 2, 21, 12 6; (b) 56, 6, 0, 12 6; (c) 55, 4, 15, 12 1E.11 Write the ground-state electron configuration for each of the following atoms: (a) sodium; (b) silicon; (c) chlorine; (d) rubidium 1E.12 Write the ground-state electron configuration for each of the following atoms: (a) titanium; (b) chromium; (c) europium; (d) krypton 1E.13 Write the ground-state electron configuration for each of the following atoms: (a) silver; (b) beryllium; (c) antimony; (d) gallium; (e) tungsten; (f) iodine 1E.14 Write the ground-state electron configuration for each of the following atoms: (a) germanium; (b) cesium; (c) iridium; (d) tellurium; (e) thallium; (f) plutonium 50 Topic 1E Many-Electron Atoms 1E.15 Which elements are predicted to have the following ground- 1E.22 How many unpaired electrons are predicted for the ground-state configuration of each of the following atoms: (a) Pb; (b) Ir; (c) Y; (d) Cd? 1E.16 Which elements are predicted to have the following ground- 1E.23 The elements Ga, Ge, As, Se, and Br lie in the same period in the periodic table Write the electron configuration expected for the ground-state atoms of these elements and predict how many unpaired electrons, if any, each atom has state electron configurations of their atoms: (a) [Kr]4d105s25p4; (b) [Ar]3d34s2; (c) [He]2s22p2; (d) [Rn]7s26d2? state electron configurations of their atoms: (a) [Ar]3d104s24p1; (b) [Ne]3s1; (c) [Kr]5s2; (d) [Xe]4f76s2? 1E.17 For each of the following ground-state atoms, predict the type of orbital (1s, 2p, 3d, 4f, etc.) from which an electron will be removed to form the 11 ion: (a) Ge; (b) Mn; (c) Ba; (d) Au 1E.18 For each of the following ground-state atoms, predict the type of orbital (1s, 2p, 3d, 4f, etc.) from which an electron will be removed to form the 11 ion: (a) Zn; (b) Cl; (c) Al; (d) Cu 1E.19 Predict the number of valence electrons present in each of the following atoms (include the outermost d-electrons): (a) N; (b) Ag; (c) Nb; (d) W 1E.20 Predict the number of valence electrons present in each of the following atoms (include the outermost d-electrons): (a) Ta; (b) Tc; (c) Te; (d) Tl 1E.21 How many unpaired electrons are predicted for the ground-state configuration of each of the following atoms: (a) Bi; (b) Si; (c) Ta; (d) Ni? 1E.24 The elements N, P, As, Sb, and Bi belong to the same group in the periodic table Write the electron configuration expected for the ground-state atoms of these elements and predict how many unpaired electrons, if any, each atom has 1E.25 Give the notation for the valence-shell configuration (including the outermost d-electrons) of (a) the alkali metals; (b)  Group 15 elements; (c) Group transition metals; (d) the “coinage” metals (Cu, Ag, Au) 1E.26 Give the notation for the valence-shell configuration (including the outermost d-electrons) of (a) the halogens; (b) the chalcogens (the Group 16 elements); (c) the transition metals in Group 5; (d) the Group 14 elements Topic 1F Periodicity 1F.1 The General Structure of the Periodic Table 1F.2 Atomic Radius 1F.3 Ionic Radius 1F.4 Ionization Energy 1F.5 Electron Affinity 1F.6 The Inert-Pair Effect 1F.7 Diagonal Relationships 1F.8 The General Properties of the Elements How is the structure of the hydrogen atom extended to other atoms? Topic 1E: Many-electron atoms How does the electronic structure of an atom relate to its position in the Periodic Table? Topic 1F: Periodicity Why Do You Need to Know This Material? The periodic table sum- The formulation of the periodic table is one of the most notable and useful achievements in chemistry because it helps to organize what would otherwise be a bewildering array of properties of the elements However, the fact that its structure corresponds to the electronic structure of atoms was unknown to its discoverers The periodic table was developed solely from a consideration of physical and chemical properties of the elements In 1869 two scientists, Lothar Meyer, a German, and Dmitri Mendeleev, a Russian (FIG 1F.1), discovered independently that the elements fell into families with similar properties when they were arranged in order of increasing atomic mass Mendeleev called this observation the periodic law One problem with Mendeleev’s table, however, was that some elements seemed to be out of place For example, when argon was isolated, it did not seem to have the correct mass for its location Its atomic weight of 40 (that is, its molar mass of 40 g?mol21) is almost the same as that of calcium, but argon is an inert gas and calcium is a reactive metal Such anomalies led scientists to question the use of atomic weight as the basis for organizing the elements In the early twentieth century, Henry Moseley examined x-ray spectra of the elements produced by bombarding a sample with a beam of electrons He realized that he could infer the atomic number itself by noting how the frequencies of the x-rays depended on the nuclear charge, and therefore on Z It was soon discovered that elements fall into the uniformly repeating pattern of the periodic table if they are organized according to atomic number rather than atomic weight marizes trends in the properties of the elements, and the ability to predict the properties of an element from its location in the periodic table is a central skill of a chemist What Do You Need to Know Already? You need to be familiar with the structure of the periodic table and its relation to the structures of many-electron atoms, as implied by the building-up principle (Topic 1E) You need to be aware of the definition of ionization energy (Topic 1D) 1F.1 The General Structure of the Periodic Table At the time the periodic table was formulated, the reason for the periodicity of the elements was a mystery Now, however, its organization is understood in terms of the electron configurations of the elements The table is divided into blocks, named for the last subshell that is occupied according to the building-up principle (s-, p-, d-, and f-blocks), as shown in Fig 1E.4, which is repeated here as FIG 1F.2 Two elements are exceptions Because it has two 1s-electrons, helium ought to lie in the s-block, but it is shown in the p-block because of its properties: it is a gas with properties closely matching those of the noble gases in Group 18, rather than the reactive metals in Group Its place in Group 18 is justified because it has a filled valence shell, like all the other Group 18 elements Hydrogen occupies a unique position in the periodic table It has one s-electron, and so it belongs in Group 1; but it is also one electron short of a noble-gas configuration, and so it can act like a member of Group 17 Because hydrogen has such a unique character, in this text it is not ascribed to any group; however, you will often see it placed in Group or Group 17, and sometimes in both THINKING POINT What are the arguments for and against including He in Group 2, above beryllium? (a) (b) FIGURE 1F.1 (a) Dmitri Ivanovitch Mendeleev (1834–1907) and (b) Lothar Meyer (1830–1895) ((a) RIA Novosti/ Science Source (b) Science Source.) 51 52 Topic 1F Periodicity FIGURE 1F.2 The names of the Period blocks of the periodic table indicate the last subshell being occupied according to the building-up principle The colors of the blocks match the colors used for the corresponding orbitals H 1s s block f block 18 p block He 13 14 15 16 17 d block [He] 2s [Ne] 3s 10 11 12 2p Ne 3p Ar [Ar] 4s 3d 4p Kr [Kr] 5s 4d 5p Xe [Xe] 6s 5d 4f 6p Rn [Rn] 7s 6d 5f 7p 14 elements The s- and p-blocks form the main groups of the periodic table The similar valenceshell electron configurations for the elements in the same main group are the reason for the similar properties of these elements The group number indicates how many valenceshell electrons are present In the s-block, the group number (1 or 2) is the same as the number of valence electrons This relation is also true for all main groups when the older practice of using Roman numerals (I–VIII) to label the groups is used However, when the modern group labels 1–18 are used, in the p-block subtract 10 from the group number to find the number of valence electrons For example, fluorine in Group 17 (old notation: Group VII) has seven valence electrons Each new period corresponds to the occupation of a shell with a higher principal quantum number This correspondence explains the different lengths of the periods: • • • • • • Period consists of only two elements, H and He, in which the single 1s-orbital of the n shell is being filled with its two electrons Period consists of the eight elements Li through Ne, in which the one 2s- and three 2p-orbitals are being filled with eight more electrons In Period (Na through Ar), the 3s- and 3p-orbitals are being occupied by eight additional electrons In Period 4, not only are the eight electrons of the 4s- and 4p-orbitals being added, so are the ten electrons of the 3d-orbitals Hence there are 18 elements in Period Period elements add another 18 electrons as the 5s-, 4d-, and 5p-orbitals are filled In Period 6, a total of 32 electrons are added, because 14 electrons are also being added to the seven 4f-orbitals The f-block elements have very similar chemical properties, because their electron configurations differ only in the population of inner f-orbitals, and electrons in these orbitals not participate much in bond formation The periodic table can be used to predict a wide range of properties, many of which are crucial for understanding chemistry The variation of effective nuclear charge, Zeff e, through the periodic table plays an important role in the explanation of periodic trends because it influences the energies and locations of the electrons in the valence shells of the atoms FIGURE 1F.3 shows the variation for the first three periods The effective charge increases from left to right across a period and falls back sharply upon going to the next period THINKING POINT Before reading further, predict how the effective nuclear charge might affect some atomic properties, such as the size of an atom or the ease with which an outer electron can be removed 1F.2 Effective nuclear charge/e, Zeff Ar P O Si Al N Mg C effective nuclear charge for the outermost valence electron with atomic number Notice that the effective nuclear charge increases from left to right across a period but drops when the outer electrons occupy a new shell (The effective nuclear charge is actually Zeffe, but Zeff itself is commonly referred to as the charge.) Be Li H 11 Atomic number, Z 13 15 17 The blocks of the periodic table are named for the last orbital to be occupied according to the building-up principle The periods are numbered according to the principal quantum number of the valence shell 1F.2 Atomic Radius Electron clouds not have sharp boundaries and so the exact radius of an atom cannot be measured However, the electron density of many-electron atoms falls off very sharply at the “edge” of the atom, and when atoms pack together in solids and bond together to form molecules, their centers are found at definite distances from one another The atomic radius of an element is defined as half the distance between the centers of neighboring atoms (1) Then: • If the element is a metal, its atomic radius is taken to be half the distance between the centers of neighboring atoms in a solid sample For instance, because the distance between neighboring nuclei in solid copper is 256 pm, the atomic radius of copper is 128 pm • If the element is a nonmetal or a metalloid, half the distance between the nuclei of atoms joined by a chemical bond is used; this radius is also called the covalent radius of the element, for reasons explained in Topic 2D For instance, the distance between the nuclei in a Cl2 molecule is 198 pm, and so the covalent radius of chlorine is 99 pm • If the element is a noble gas, the van der Waals radius is used, which is half the distance between the centers of neighboring atoms in a sample of the solidified gas The atomic radii of the noble gases listed in Appendix 2D are all van der Waals radii Because the atoms in a sample of a noble gas are not chemically bonded together, van der Waals radii are generally much larger than covalent radii and are best not included in the discussion of trends FIGURE 1F.4 shows the atomic radii of some main-group elements and FIG 1F.5 shows the variation in atomic radius with atomic number Note the periodic, sawtooth pattern A feature to note is that: • Atomic radius generally decreases from left to right across a period and increases down a group The increase in atomic radius down a group, such as that from Li to Cs, makes sense: with each new period, the outermost electrons occupy shells with increasing principal quantum number and therefore lie farther from the nucleus The decrease across a period, such as that from Li to Ne, might be surprising at first sight because the number of 53 FIGURE 1F.3 The variation of the Na B He S F Cl Ne Atomic Radius 2r Atomic radius 54 Topic 1F Periodicity FIGURE 1F.4 The atomic radii (in picometers) of the main-group elements The radii decrease from left to right in a period and increase down a group Atomic radii, including those of the d-block elements, are listed in Appendix 2D FIGURE 1F.5 The periodic variation in 13 14 15 16 17 18 Li 152 Be 113 B 83 C 77 N 75 O 66 F 71 Ne Na 154 Mg 160 Al 143 Si 117 P 110 S 104 Cl 99 Ar K 227 Ca 197 Ga 122 Ge 122 As 121 Se 117 Br 114 Kr Rb 248 Sr 215 In 163 Sn 141 Sb 141 Te 137 I 133 Xe Cs 265 Ba 217 Tl 170 Pb 175 Bi 155 Po 167 At Rn 300 Cs K Atomic radius, r/pm the atomic radii of the elements The decrease across a period can be explained in terms of the effect of increasing effective nuclear charge and the increase down a group can be explained by the occupation of shells with increasing principal quantum number Rb Na 200 Po Li I Cl 100 Br F s block p block d block f block Atomic number, Z electrons is increasing along with the number of protons The explanation is that the new electrons are in the same shell of the atom and about as close to the nucleus as other electrons in the same shell However, because they are spread out in the shell, the electrons not shield one another well from the nuclear charge; so the effective nuclear charge increases across the period The increasing effective nuclear charge draws the electrons in As a result, the atom is more compact and there is a trend for atomic radii to increase diagonally from the upper right of the periodic table to the lower left THINKING POINT Which currently known element has the biggest atoms? Atomic radii generally decrease from left to right across a period as the effective atomic number increases, and they increase down a group as successive shells are occupied 1F.3 Ionic Radius ranion + rcation – Ionic radius + The radii of ions differ markedly from the radii of their parent atoms As described in Fundamentals C, each ion in an ionic solid is surrounded by ions with the opposite charge The ionic radius of an element is its share of the distance between neighboring ions in an ionic solid (2) The distance between the centers of a neighboring cation and anion is the sum of the two ionic radii In practice, the radius of the oxide ion is taken to be 140 pm and the radii of other ions are calculated on the basis of that value For example, because the distance between the centers of neighboring Mg21 and O22 ions in magnesium oxide is 212 pm, the radius of the Mg21 ion is reported as 212 pm 140 pm 72 pm 1F.3 Li+ 76 Na+ 102 K+ 138 Be 13 B 2+ 45 Mg2+ Rb+ 152 Sr2+ 118 Cs+ 167 Ba2+ 135 15 16 17 18 C N3– 171 O2– 140 F– 133 Ne 23 Si P3– 212 S2– 184 Cl– 181 Ar Ge As3– 222 Se2– 198 Br– 196 Kr Sn Sb Te2– 221 I– 220 Xe Al3+ 54 72 Ca2+ 100 14 3+ Ga3+ 62 In3+ 80 Tl3+ 89 Pb Bi Po At Li+ Li Be2+ Na+ Na Mg2+ O O2– F S S2– Cl cations are typically smaller than their parent atoms, whereas anions are larger—in some cases, very much larger FIGURE 1F.6 illustrates the trends in ionic radii, and FIG 1F.7 shows the relative sizes of some ions and their parent atoms All cations are smaller than their parent atoms, because the atom loses its valence electrons to form the cation and exposes its core, which is generally much smaller than the parent atom For example, the atomic radius of Li, with the configuration 1s22s1, is 152 pm, but the ionic radius of Li1, the bare heliumlike 1s2 core of the parent atom, is only 76 pm This size difference is comparable to that between a cherry and its pit Atoms in the same main group tend to form ions with the same charge Like atomic radii, the radii of these ions increase down each group because the core electrons occupy shells with higher principal quantum numbers Figure 1F.7 shows that anions are larger than their parent atoms The reason can be traced to the increased number of electrons in the valence shell of the anion and the repulsive effects exerted by electrons on one another The variation in radii of anions shows the same diagonal trend as that for atoms and cations, with the smallest at the upper right of the periodic table, close to fluorine: F– Cl– some cations and anions compared with their parent atoms Note that cations (pink) are smaller than their parent atoms (gray), whereas anions (green) are larger Atoms and ions with the same number of electrons are called isoelectronic For example, Na1, F2, and Mg21 are isoelectronic All three ions have the same electron configuration, [He]2s22p6, but their radii differ because they have different nuclear charges (see Fig 1F.3) The Mg21 ion has the largest nuclear charge; so it has the strongest attraction for the electrons and therefore the smallest radius The F2 ion has the lowest nuclear charge of the three isoelectronic ions and, as a result, it has the largest radius EXAMPLE 1F.1 Deciding the relative sizes of ions Mineralogists and geologists often need to identify the relative sizes of atoms to judge whether one mineral might be modified by the inclusion of “alien” ions For example, the different colors of some gemstones result from this type of insertion Arrange each of the following pairs of ions in order of increasing ionic radius: (a) Mg21 and Ca21; (b) O22 and F2 PLAN The smaller member of a pair of isoelectronic ions in the same period will be an ion of an element that lies farther to the right in a period, because that ion has the greater effective nuclear charge If the two ions are in the same group, the smaller ion will be the one that lies higher in the group, because its outermost electrons are closer to the nucleus SOLVE 2+ Mg21 has the smaller ionic radius Mg FIGURE 1F.7 The relative sizes of Cations are smaller than their parent atoms, whereas anions are larger (a) Mg lies above Ca in Group Be Rn FIGURE 1F.6 The ionic radii (in picometers) of the ions of the main-group elements Note that • 55 Ionic Radius Mg 72 pm 2+ Ca 100 pm 56 Topic 1F Periodicity (b) F lies to the right of O in Period F2 has the smaller ionic radius 14 15 16 17 140 pm 2– – O F 133 pm EVALUATE Appendix 2C shows that the actual values are (a) 72 pm for Mg 21 and 100 pm for Ca21; (b) 133 pm for F2 and 140 pm for O22 Self-test 1F.1A Arrange each of the following pairs of ions in order of increasing ionic radius: (a) Mg21 and Al31; (b) O22 and S22 [Answer: (a) r(Al31) , r(Mg21); (b) r(O22) , r(S22)] Self-test 1F.1B Arrange each of the following pairs of ions in order of increasing ionic radius: (a) Ca21 and K1; (b) S22 and Cl2 Related Exercises 1F.3, 1F.4 Ionic radii generally increase down a group and decrease from left to right across a period Cations are smaller than their parent atoms and anions are larger 1F.4 Ionization Energy By referring to the minimum energy you don’t have to worry about the kinetic energy of the electron: it is assumed to be stationary Ionization can be achieved by using a higher energy, but then the electron would carry away the excess energy as kinetic energy As explained in Topic 2A, the formation of a bond in an ionic compound depends on the removal of one or more electrons from one atom and the transfer of those electrons to another atom The energy needed to remove electrons from atoms is therefore of central importance for understanding their chemical properties As remarked in Topic 1D, the ionization energy, I, is the minimum energy needed to remove an electron from an atom in the gas phase Specifically: J(g) ¡ J (g) e (g) I E(J ) E(J) (1) where E(J) is the energy of species J Ionization energies are reported either as molar quantities in kilojoules per mole (kJ?mol21) or in electronvolts (eV), the change in energy of an electron when it moves through a potential difference of volt (1 eV 1.602 10219 J) The first ionization energy, I1, is the minimum energy needed to remove an electron from a neutral atom in the gas phase For example, for copper, Cu(g) ¡ Cu (g) e (g) energy required I1 (7.73 eV, 746 kJ?mol 21 ) The second ionization energy, I2, of an element is the minimum energy needed to remove an electron from a singly charged gas-phase cation For copper, Cu (g) ¡ Cu21 (g) e (g) energy required I2 (20.29 eV, 1958 kJ?mol 21 ) Because ionization energy is a measure of how difficult it is to remove an electron, elements with low ionization energies can be expected to form cations readily and to conduct electricity (which requires that some electrons be free to move) in their solid (and liquid) forms Elements with high ionization energies are unlikely to form cations and are unlikely to conduct electricity THINKING POINT Why is the second ionization energy of an atom always higher than its first ionization energy? As you can see from FIG 1F.8: • • First ionization energies typically decrease down a group First ionization energies generally increase across a period 1F.4 H 1310 18 13 14 15 16 17 He 2370 18/VIII Li 519 Be 900 B 799 C 1090 N 1400 O 1310 F 1680 Ne 2080 Na Mg Al Si P S Cl Ar 494 736 577 786 1011 1000 1255 1520 K 418 Ca 590 Ga 577 Ge 784 As 947 Se 941 Br 1140 Kr 1350 Rb 402 Sr 548 In 556 Sn 707 Sb 834 Te 870 I 1008 Xe 1170 Cs 376 Ba 502 Tl 590 Pb 716 Bi 703 Po 812 At 1037 Rn 1036 FIGURE 1F.8 The first ionization energies of the main-group elements, in kilojoules per mole In general, low values are found at the lower left of the table and high values are found at the upper right The decrease down a group can be explained by the finding that, in successive periods, the outermost electron occupies a shell that is farther from the nucleus and is therefore less tightly bound Therefore, it takes less energy to remove an electron from a cesium atom, for instance, than from a sodium atom With few exceptions, the first ionization energy rises from left to right across a period (FIG 1F.9) This trend can be traced to the increase in effective nuclear charge across a period The small departures from this trend arise from repulsions between electrons, particularly electrons occupying the same orbital For example, the ionization energy of oxygen is slightly lower than that of nitrogen because in a nitrogen atom each p-orbital has one electron, but in oxygen the eighth electron is paired with an electron already occupying an orbital The repulsion between the two electrons in the same orbital raises their energy and makes one of them easier to remove from the atom than if the two electrons had been in different orbitals FIGURE 1F.10 shows that the second ionization energy of an element is always higher than its first ionization energy It takes more energy to remove an electron from a positively charged ion than from a neutral atom For the Group elements, the second ionization energy is considerably larger than the first; in Group 2, however, the two ionization energies have similar values This difference makes sense, because the Group elements have an ns1 valence-shell electron configuration Although the removal of the first electron 14 800 25 000 First Second Third Fourth He s block p block d block f block Ionization energy, I Ne 7300 3660 Ar 4560 3070 Kr H 57 Ionization Energy Xe 1760 Rn 2420 519 900 799 494 418 Li Be Na K B FIGURE 1F.10 The successive Li Na K Rb Cs Fr Atomic number, Z FIGURE 1F.9 The periodic variation of the first ionization energies of the elements ionization energies of a selection of main-group elements Note the great increase in energy required to remove an electron from an inner shell In each case, the blue outline denotes ionization from the valence shell 58 Topic 1F Periodicity Metal block requires only a small input of energy, the second electron must come from the noble-gas core The core electrons have lower principal quantum numbers and are much closer to the nucleus They are strongly attracted to it and a lot of energy is needed to remove them Self-test 1F.2A Account for the slight decrease in first ionization energy between beryllium and boron Cation Electron sea FIGURE 1F.11 A block of metal consists of an array of cations (the spheres) surrounded by a sea of electrons The charge of the electron sea cancels the charges of the cations The electrons of the sea are mobile and can move past the cations quite easily and hence conduct an electric current ANIMATION FIGURE 1F.11 [Answer: Boron loses an electron more easily from a higher-energy subshell than beryllium does.] Self-test 1F.2B Account for the large decrease in third ionization energy between beryllium and boron The low ionization energies of elements at the lower left of the periodic table account for their metallic character A block of metal consists of a collection of cations of the element surrounded by a sea of valence electrons that the atoms have lost (FIG 1F.11) Only elements with low ionization energies—the members of the s-block, the d-block, the f-block, and the lower left of the p-block—can form metallic solids, because only they can lose electrons easily The elements at the upper right of the periodic table have high ionization energies, so they not readily lose electrons and are therefore not metals Note that your knowledge of electronic structure has helped you to understand a major feature of the periodic table—in this case, why the metals are found toward the lower left and the nonmetals are found toward the upper right The first ionization energy is highest for elements close to helium and is lowest for elements close to cesium Second ionization energies are higher than first ionization energies (of the same element) and very much higher if the electron is to be removed from a closed shell Metals are found toward the lower left of the periodic table because these elements have low ionization energies and can readily lose their electrons 1F.5 Electron Affinity To predict some chemical properties, you need to know how the energy changes when an electron attaches to an atom The electron affinity, Eea, of an element is the energy released when an electron is added to a gas-phase atom A positive electron affinity means that energy is released when an electron attaches to an atom A negative electron affinity means that energy must be supplied to push an electron onto an atom This convention matches the everyday meaning of the term “affinity.” More formally, the electron affinity of an element X is defined as X(g) e (g) ¡ X (g) Eea (X) E(X) E(X ) (2) where E(X) is the energy of a gas-phase X atom and E(X ) is the energy of the gas-phase anion For instance, the electron affinity of chlorine is the energy released in the process Cl(g) e (g) ¡ Cl (g) In some books, you will see electron affinity defined with an opposite-sign convention Those values are actually the electron-gain enthalpies (Topic 4C) energy released Eea (3.62 eV, 349 kJ?mol 21 ) Because the electron has a lower energy when it occupies one of the atom’s orbitals, the difference E(Cl) E(Cl2) is positive and the electron affinity of chlorine is positive Like ionization energies, electron affinities are reported either in electronvolts for a single atom or in joules per mole of atoms FIGURE 1F.12 shows the variation in electron affinity in the main groups of the periodic table It is much less periodic than variations in radius and ionization energy However, one broad trend is clearly visible With the exception of the noble gases: • Electron affinities are highest toward the right of the periodic table This trend is particularly true in the upper right, close to oxygen, sulfur, and the halogens In these atoms, the incoming electron occupies a p-orbital close to a nucleus with a high effective charge and can experience its attraction quite strongly The noble gases have negative electron affinities because any electron added to them must occupy an orbital outside a closed shell and far from the nucleus: this process requires energy, and so the electron affinity is negative 1F.5 18 Na Mg Al Si P +53 ≤0 +43 +134 +72 16 O +141 –844 S +200, –532 K +48 Ca +2 Ga +29 Ge +116 As +78 Rb +47 Sr +5 In +29 Sn +116 Cs +46 Ba +14 Tl +19 Pb +35 13 14 15 Li +60 Be ≤0 B +27 C +122 N –7 17 He