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Section 3 5b, 5C Main Ideas Atomic radii are related to electron configuration. Removing electrons from atoms to form ions requires energy. Adding electrons to atoms to form ions also involves energy. When atoms become ions, their radii change. Only the outer electrons are involved in forming compounds. Atoms have different abilities to capture electrons. The properties of d-block metals do not vary much. 5B use the Periodic Table to identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals; 5C use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity, and ionization energy Electron Configuration and Periodic Properties Key Terms atomic radius ion ionization ionization energy electron affinity cation anion valence electron electronegativity So far, you have learned that the elements are arranged in the periodic table according to their atomic number and that there is a rough correlation between the arrangement of the elements and their electron configurations. In this section, the relationship between the periodic law and electron configurations will be further explored. Main Idea 5C Atomic radii are related to electron configuration. Ideally, the size of an atom is defined by the edge of its largest orbital. However, this boundary is fuzzy and varies under different conditions. Therefore, the conditions under which the atom exists must be specified to estimate its size. One way to express an atom’s radius is to measure the distance between the nuclei of two identical atoms that are chemically bonded together and then divide this distance by two. Atomic radius may be defined as one-h alf the distance between the nuclei of identical atoms that are bonded together. This can be seen in Figure 3.1 on the next page. Period Trends Figure 3.2 gives the atomic radii of the elements, and Figure 3.3 (on the next Go online Chemistry HMDScience.com Period Trends (Interaction) spread) presents this information graphically. Note that there is a gradual decrease in atomic radii across the second period from lithium, Li, to neon, Ne. The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus. As electrons add to s and p sublevels in the same main energy level, they are gradually pulled closer to the more highly charged nucleus. This increased pull results in a decrease in atomic radii. The attraction of the nucleus is somewhat offset by repulsion among the increased number of electrons in the same outer energy level. As a result, the difference in radii between neighboring atoms in each period grows smaller, as shown in Figure 3.2. Group Trends Examine the atomic radii of the Group 1 elements in Figure 3.2. Notice that the radii of the elements increase as you read down the group. 148 Chapter 5 Figure 3.1 Atomic Radii One method of Chlorine nucleus determining atomic radius is to measure the distance between the nuclei of two identical atoms that are bonded together in an element or compound and then divide this distance by two. The atomic radius of a chlorine atom, for example, is about 100 picometers (pm). Atomic radius 99 pm 198 pm Chlorine nucleus Distance between nuclei CHECK FOR UNDERSTANDING As electrons occupy sublevels in successively higher mainenergy levels farther from the nucleus, the sizes of the atoms increase. In general, the atomic radii of the main-g roup elements increase down a group. Explain Why is it necessary for the radii of atoms to be expressed as the distance between the nuclei of two identical atoms bonded together? (Refer to Figure 3.1.) Now examine the radii of the Group 13 elements. Although gallium, Ga, follows aluminum, Al, it has a slightly smaller atomic radius than does aluminum. This is because gallium, unlike aluminum, is preceded in its period by the 10 d-block elements. The expected increase in gallium’s radius caused by the filling of the fourth main-energy level is outweighed by a shrinking of the electron cloud caused by a nuclear charge that is considerably higher than that of a luminum. Figure 3.2 Decreasing Radii Atomic radii decrease from left to right across a period and increase down a group. Values are given in picometers (pm). 1 2 H 37 Group 1 Group 2 Li 3 Be 4 152 3 Period 31 Group 13 Group 14 Group 15 Group 16 Group 17 B 5 C 6 N 7 O 8 F 9 Ne 10 Relative atomic size 77 85 Cs Fr 13 Si 75 14 P 73 15 S 72 16 Cl 2 71 17 Ar 18 118 110 103 100 98 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 143 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 197 180 134 128 127 40 Nb 41 Mo 42 Tc 160 57 Hf 183 88 Ac 220 147 39 Zr 56 La 222 87 Ra 162 38 Y 215 55 Ba 270 Al 77 1 3 160 Rb 37 Sr 265 7 Group 18 He 2 Atomic radius 112 19 Ca 248 6 6 146 72 Ta 159 139 73 W 146 125 124 43 Ru 44 Rh 45 Pd 136 74 Re 139 126 134 75 Os 137 134 76 Ir 135 136 — — — — — 135 144 149 49 Sn 167 78 Au 79 Hg 80 Tl 122 120 50 Sb 140 81 Pb 119 51 Te 140 82 Bi 142 83 Po — — — 133 5 131 85 Rn 86 140 4 112 53 Xe 54 84 At 144 151 170 175 150 168 110 Rg 111 Cn 112 113 Fl 114 115 Lv 116 — — 114 52 I 139 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 188 134 46 Ag 47 Cd 48 In 137 77 Pt 128 Period K 227 5 C Atomic number Na 11 Mg 12 186 4 Atomic symbol 1 6 141 117 118 — 7 Lanthanide series Ce 58 Pr 182 Th 182 90 Pa 179 59 Nd 60 Pm 61 Sm 62 Eu 91 U 163 181 183 180 92 Np 93 Pu 156 155 63 Gd 64 Tb 208 180 65 Dy 66 Ho 67 Er 177 94 Am 95 Cm 96 Bk 159 173 174 178 97 Cf — 176 98 Es 186 186 68 Tm 69 Yb 70 Lu 176 176 — 71 174 99 Fm 100 Md101 No 102 Lr 103 — — — — Actinide series The Periodic Law 149 Figure 3.3 Periodic Trends The plot of atomic radius versus atomic number shows period and group trends. Atomic Radius vs. Atomic Number 300 Period 2 Period 3 Period 4 Period 5 Period 6 Cs Fr Period 7 Rb 250 K Atomic radius (pm) 200 Es Na Li 150 Rn Xe Kr 100 Ar Ne 50 H He 0 0 10 20 30 40 50 60 70 80 90 100 Atomic number 5C Go Online Learn It! Video Atomic Radius HMDScience.com Sample Problem E Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus, P, which has the largest atomic radius? Explain your answer in terms of trends in the periodic t able. Solve All of the elements are in the third period. Of the four, sodium has the lowest atomic number and is the first element in the period. This means sodium has the smallest nucleus and, therefore, the least attraction on the electrons. Therefore, sodium has the largest atomic radius, because atomic radii decrease across a period. Answers in Appendix E 1. Which of the following elements has the largest atomic radius: Li, O, C, or F? Which has the smallest atomic radius? 2. Of the elements calcium, Ca, beryllium, Be, barium, Ba, and strontium, Sr, which has the largest atomic radius? Explain your answer in terms of trends in the periodic t able. 3. Of the elements aluminum, Al, magnesium, Mg, silicon, Si, and sodium, Na, which has the smallest atomic radius? Explain your answer in terms of trends in the periodic table. 150 Chapter 5 5B, 5C Main Idea Removing electrons from atoms to form ions requires energy. An electron can be removed from an atom if enough energy is supplied. Using A as a symbol for an atom of any element, the process can be expressed as follows. A + energy → A+ + eThe A+ represents an ion of element A with a single positive charge, referred to as a 1+ ion. An ion is an atom or group of bonded atoms that has a positive or negative charge. Sodium, for example, forms an Na+ ion. Any process that results in the formation of an ion is referred to as ionization. To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies. The energy required to remove one electron from a neutral atom of an element is the ionization energy, IE (or first ionization energy, IE1). To avoid the influence of nearby atoms, measurements of ionization energies are made on isolated atoms in the gas phase. Figure 3.4 gives the first ionization energies for the elements in kilojoules per mole (kJ/mol). Figure 3.5, on the next page, presents this information graphically. Figure 3.4 Ionization Energy In general, first ionization energies increase across a period and decrease down a group. Values are given in kJ/mol. 1 2 Li 520 11 5 6 7 Be 12 Mg 19 20 738 2 Group 13 Group 14 Group 15 Group 16 Group 17 Symbol 1086 900 Na 496 C 4 Group 18 5 B First ionization energy 21 22 24 25 26 27 28 29 30 2372 10 1402 1314 1681 F Ne 13 14 15 16 17 18 578 787 1012 1000 1251 1521 31 32 33 34 35 36 P O 9 1086 Si N 8 S Cl Sc 633 659 651 V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se 941 1140 1351 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 653 717 762 760 737 746 906 579 762 947 Sr 550 600 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te 869 1008 I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 640 652 684 702 710 720 804 731 868 558 709 834 1170 Cs Ba La Hf Ta W Re Os 839 878 Ir Pt Au Hg 1007 589 Tl Pb Bi Po At Rn 87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 376 Fr — 503 Ra 509 538 Ac 490 659 Rf — 761 Db — 770 Sg — 760 Bh — Hs — Mt — 868 Ds — 890 Rg — Cn — 716 Fl 703 — 812 Lv — 3 4 Kr Rb 403 2 Ar Ca 590 Br 2081 K 419 Ti 23 C 7 He 801 Al Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 6 1 1038 5 Period Period 4 6 1312 Group 1 Group 2 3 3 Atomic number 1 H 6 7 — Lanthanide series 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 90 91 92 93 94 95 96 97 98 99 100 101 102 103 578 581 642 — 534 Th 527 Pa 587 570 Actinide series 533 U 598 536 Np 600 545 Pu 585 547 592 Am Cm 566 Bk 601 573 Cf 608 581 Es 619 589 Fm 627 597 Md 635 603 No 523 Lr The Periodic Law 151 Figure 3.5 Ionization Energy Plot of first ionization energy, IE 1, versus atomic number. As atomic number increases, both the period and the group trends become less pronounced. First Ionization Energy vs. Atomic Number 3000 Period 2 2500 Period 3 Period 4 Period 5 Period 6 Period 7 He First ionization energy (kJ/mol) Ne 2000 Ar 1500 Kr H Xe Rn 1000 No Li 500 0 Ra 0 Na 10 K 20 Rb 30 40 Cs 50 60 Atomic number 70 80 90 100 Period Trends GO online Chemistry HMDScience.com mc06sec05000060a Period Trends 1st Pass 6/18/04 Nanda Patel In Figures 3.4 and 3.5, examine the ionization energies for the first and last elements in each period. You can see that the Group 1 metals have the lowest first ionization energies in their respective periods. Therefore, they lose electrons most easily. This ease of electron loss is a major reason for the high reactivity of the Group 1 (alkali) metals. The Group 18 elements, the noble gases, have the highest ionization energies. They do not lose electrons easily. The low reactivity of the noble gases is partly based on this difficulty of electron removal. In general, ionization energies of the main-group elements increase across each period. This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level. Increasing nuclear charge is responsible for both increasing ionization energy and decreasing radii across the periods. Note that, in general, nonmetals have higher ionization energies than metals do. In each period, the element of Group 1 has the lowest ionization energy, and the element of Group 18 has the highest ionization energy. 152 Chapter 5 Group Trends Among the main-g roup elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus. Therefore, they are removed more easily. Also, as atomic number increases going down a group, more electrons lie between the nucleus and the electrons in the highest occupied energy levels. This partially shields the outer electrons from the effect of the nuclear charge. Together, these influences overcome the attraction of the electrons due to the increasing nuclear c harge. Removing Electrons from Positive Ions With sufficient energy, electrons can be removed from positive ions as well as from neutral atoms. The energies for removal of additional electrons from an atom are referred to as the second ionization energy (IE2), third ionization energy (IE3), and so on. Figure 3.6 shows the first five ionization energies for the elements of the first, second, and third periods. You can see that the second ionization energy is always higher than the first, the third is always higher than the second, and so on. Figure 3.6 Ionization Energies (in kj/mol) for Elements of Periods 1–3 Period 1 H IE1 IE2 IE3 1312 He Period 2 Li Be B C N O F Ne 2372 520 900 801 1086 1402 1314 1681 2081 5250 7298 1757 2427 2353 2856 3388 3374 3952 11 815 14 849 3660 4621 4578 5300 6050 6122 21 007 25 026 6223 7475 7469 8408 9370 32 827 37 830 9445 10 990 11 023 12 178 P S Cl Ar IE4 IE5 Period 3 Na Mg Al Si IE1 496 738 578 787 1012 1000 1251 1521 IE2 4562 1451 1817 1577 1903 2251 2297 2666 IE3 6912 7733 2745 3232 2912 3361 3822 3931 IE4 9544 10 540 11 578 4356 4957 4564 5158 5771 IE5 13 353 13 628 14 831 16 091 6274 7013 6540 7238 The Periodic Law 153 CHECK FOR UNDERSTANDING Explain Explain in your own words why ionization energies increase as successive electrons are removed from an ion. This rise in ionization energies is understandable when you consider what is occurring. The first electron is removed from a neutral atom. However, the second electron is removed from a positive ion, so it takes more energy to remove it. Therefore, it is harder to remove each additional electron because it is being removed from an even more positive ion. The electrons in the outermost energy level are called valence electrons. Those electrons in the lower levels are called core electrons. Once all of the valence electrons have been removed, it requires a lot more energy to remove successive core electrons. Look at the pattern in Figure 3.6 again. The green boxes are the IE values to remove the first core electron, after all the valence electrons have been removed. Using Na (sodium) with the electron configuration 1s22p23s1 as an example, the 1s, 2s, and 2p electrons are the core electrons, and the 3s electron is the valence electron. Notice that it takes a lot more energy to remove the 2p electron from the core than it does the 3s valence electron. Also consider Mg (magnesium). Magnesium’s third ionization energy is much greater than its first or second ionization energies. This is because Mg has two valance electrons, and so the third ionization energy is the energy required to remove a core electron. This pattern continues for the remainder of the elements in the third period. Periodic Trends in Ionization Energy Sample Problem F Consider two main-group elements, A and B. Element A has a first ionization energy of 419 kJ/mol. Element B has a first ionization energy of 1000 kJ/mol. Decide if each element is more likely to be in the s-block or p-block. Which element is more likely to form a positive i on? SOLVE Element A has a lower ionization energy, which means that atoms of A lose electrons easily. Therefore, element A is most likely to be an s-block metal, because ionization energies increase across the periods. Element B has a higher ionization energy, which means that atoms of B have difficulty losing electrons. Element B would most likely lie at the end of a period in the p-block. Element A is more likely to form a positive ion, because it has a much lower ionization energy than element B does. Answers in Appendix E 1. Consider four hypothetical main-group e lements, Q, R, T, and X, that have the outer electron configurations indicated below. Then, answer the questions that follow. Q: 3s23p5 R: 3s1 T: 4d105s25p5 X: 4d105s25p1 a. Identify the block location of each hypothetical main-g roup element. b. Which of these elements are in the same period? Which are in the same g roup? c. Which element would you expect to have the highest first ionization energy? Which would have the lowest first ionization e nergy? d. Which element would you expect to have the highest second ionization e nergy? e. Which of the elements is most likely to form a 1+ ion? 154 Chapter 5 5B, 5C Main Idea Adding electrons to atoms to form ions also involves energy. Neutral atoms can also acquire electrons. The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. Most atoms release energy when they acquire an e lectron. A + e- → A- + energy On the other hand, some atoms must be “forced” to gain an electron by the addition of energy. A + e- + energy → AThe quantity of energy absorbed would be represented by a positive number, but ions produced in this way are very unstable, and hence the electron affinity for them is very difficult to determine. An ion produced in this way will be unstable and will lose the added electron spontaneously. Figure 3.7 shows the electron affinity in kilojoules per mole for the elements. Positive electron affinities, because they are so difficult to determine with any accuracy, are denoted in Figure 3.7 by “(0).” Figure 3.8, on the next page, presents these data g raphically. Figure 3.7 Period Trends Periodic Table of Electron Affinities (kJ/mol) The values Among the elements of each period, the halogens (Group 17) gain electrons most readily. This is shown in Figure 3.7 by the large negative values of halogens’ electron affinities and is a major reason for the high reactivity levels of Group 17 elements. In general, as electrons add to the same p sublevel of atoms with increasing nuclear charge, electron affinities become more negative across each period within the p-block. 1 2 3 5 6 7 6 –75.4 Group 1 Group 2 3 Li Be 11 12 –61.8 Mg 19 20 (0) 21 Sc 22 Ca (0) –18.8 –7.9 37 38 39 40 Rb Sr (0) –30.7 55 56 57 –48.6 Y Ti Zr –52.5 Cr –66.6 Mn Co (0) –16.3 –66.1 41 42 43 44 45 –105 –113.7 74 75 76 77 (0) — Fe 27 73 104 (0) 26 –89.3 89 –47.0 25 72 88 Ac V 24 –42.6 87 –50 23 Rf — Ta –74.6 –55 –32.2 W –81.5 Re –15 –110 –156.5 105 106 107 108 109 Db — Sg — Bh — Os Rh Hs — Ir Mt — 28 Ni 29 30 6 7 –27.7 C –126.3 N (0) 13 14 15 –44.1 –138.5 –74.6 31 32 33 Al Ru Hf Ra Electron affinity Tc La Fr B Mo Ba (0) 5 Nb Cs –47.2 Symbol Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 K –50.1 2 Group 13 Group 14 Group 15 Group 16 Group 17 –126.3 (0) Na –54.8 C 4 Group 18 Si P –115.6 Cu –122.8 Zn Ga Ge As –81 46 47 48 49 50 51 Pd (0) –30 –135 –55.7 Ag –130.2 Cd In –30 –120 –107 78 79 80 81 82 83 Pt (0) –212.8 Au –230.9 Hg 110 111 Ds — Rg — Sn Sb (0) Tl –20 Pb –36 –94.6 112 113 114 115 Cn — Fl — Bi 8 O 9 He (0) 10 –146.1 F –339.9 Ne 16 17 18 S (0) –207.7 Cl –361.7 Ar 34 35 36 Se (0) –202.1 Br –336.5 Kr 52 53 54 Te (0) –197.1 I –305.9 Xe 84 85 86 Po (0) –190 At –280 Rn 116 117 118 Lv 1 (0) 2 3 4 5 Period Period 4 Atomic number 1 H listed in parentheses in this periodic table of electron affinities are approximate. Electron affinity is estimated to be -50 kJ/mol for each of the lanthanides and 0 kJ/mol for each of the actinides. 6 7 — The Periodic Law 155 Electron Affinity vs. Atomic Number 0 Electron affinity (kJ/mol) –100 Ne He H Li Ar Na Kr Xe Ra Fr Cs Rb K Rn –200 –300 –400 Period 2 0 Period 3 10 20 Period 4 30 Period 5 40 50 60 Period 6 70 80 Period 7 90 100 Atomic number Figure 3.8 Electron Affinity and Atomic Numbers The plot of electron affinity versus atomic number shows that most atoms release energy when they acquire an electron, as indicated by negative values. An exception to this trend occurs between Groups 14 and 15. Compare the electron affinities of carbon ([He]2s22p2) and nitrogen ([He]2s22p3). Adding an electron to a carbon atom gives a half-filled p sublevel. This occurs more easily than forcing an electron to pair with another electron in an orbital of the already half-filled p sublevel of a nitrogen atom. Group Trends Trends for electron affinities within groups are not as regular as trends for ionization energies. As a general rule, electrons are added with greater difficulty down a group. This pattern is a result of two competing factors. The first is an increase in nuclear charge down a group, which i ncreases electron affinities. The second is an increase in atomic radius down a group, which decreases electron affinities. In general, the size effect predominates. But there are exceptions, especially among the heavy transition metals, which tend to be the same size or even decrease in radius down a group. Adding Electrons to Negative Ions For an isolated ion in the gas phase, it is more difficult to add a second electron to an already negatively charged ion. Therefore, second electron affinities are all positive. Certain p-block nonmetals tend to form negative ions that have nobleg as configurations. The halogens do so by adding one electron. For example, chlorine has the configuration [Ne]3s23p5. 156 Chapter 5 An atom of chlorine achieves the configuration of the noble gas argon by adding an electron to form the ion Cl- ([Ne]3s23p6). Adding another electron is so difficult that Cl2- never occurs. Atoms of Group 16 elements are present in many compounds as 2- ions. For example, oxygen ([He]2s22p4) achieves the configuration of the noble gas neon by adding two electrons to form the ion O2-([He]2s22p6). 5C Main Idea When atoms become ions, their radii change. Figure 3.9 shows the radii of some of the most common ions of the elements. Positive and negative ions have specific names. A positive ion is known as a cation. The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius, because the removal of the highest-energy-level electrons results in a smaller electron cloud. Also, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge. A negative ion is known as an anion. The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. This is because the total positive charge of the nucleus remains unchanged when an electron is added to an atom or an ion. So the electrons are not drawn to the nucleus as strongly as they were before the addition of the extra electron. The electron cloud also spreads out because of greater repulsion between the increased number of electrons. Figure 3.9 Period Trends Periodic Table of Ionic Radii (pm) The ionic radii of the ions most Within each period of the periodic table, the metals at the left tend to form cations, and the nonmetals at the right tend to form anions. Cationic radii decrease across a period, because the electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level. Atomic number 1 H– 154 Relative atomic size Group 1 Group 2 2 4 Li+ 76 Be2+ 45 C Ionic symbol 2 4– 260 Relative ionic size Group 13 Group 14 Group 15 Group 16 Group 17 He Ionic radius (pm) B 5 11 12 13 Na+ Mg2+ Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 Al3+ 21 31 102 19 72 20 22 23 24 25 26 27 28 29 30 6 — C4– 7 260 N3– 8 146 O2– 9 10 140 F– 133 Ne 2 14 15 16 17 18 54 Si P3– S2– Cl– 181 Ar — 32 33 62 Ge — As 50 51 212 184 34 1 — — 3 35 36 — Se2– 198 Br– 196 Kr — 4 4 K+ 138 Ca2+ 100 Sc3+ 37 38 39 75 Ti2+ 86 V2+ 40 41 90 Zr — Nb 79 Cr2+ 42 80 Mn2+ 83 Fe2+ 43 44 — Tc — Ru 78 Co2+ 45 65 Ni2+ 46 69 Cu2+ 47 73 Zn2+ 74 Ga3+ 48 49 52 53 Period Period 3 3 Group 18 6 1 common in chemical compounds are shown. Cations are smaller and anions are larger than the atoms from which they are formed. — 54 5 5 Rb+ 152 Sr2+ 118 Y3+ 55 56 57 — Mo 72 73 74 75 76 167 Ba2+ 136 La3+ 116 Hf — Ta — W — Re — Os — Rh3+ 77 67 Pd2+ 78 86 Ag+ 115 Cd2+ 79 95 In3+ 80 81 80 Sn2+ 118 Sb3+ 82 83 76 Te2– 221 I– 220 Xe 84 85 86 89 Pb2+ 119 Bi3+ 103 Po — At — Rn — 6 6 Cs+ 87 88 89 Fr+ Ra2+ Ac3+ 104 105 106 107 108 — Ir 109 — Pt2+ 110 80 Au+ 137 Hg2+ 119 Tl3+ 111 112 113 114 115 117 116 — 118 7 7 180 148 111 Rf — Db — Sg — Bh — Hs — Mt — Ds — Rg — Cn — Fl — Lv — The Periodic Law 157 Starting with Group 15, in which atoms assume stable noble-gas configurations by gaining three electrons, anions are more common than cations. Anionic radii decrease across each period for the elements in Groups 15–18. The reasons for this trend are the same as the reasons that cationic radii decrease from left to right across a period. Group Trends The outer electrons in both cations and anions are in higher energy levels as one reads down a group. Thus, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii. Main Idea Only the outer electrons are involved in forming compounds. Chemical compounds form because electrons are lost, gained, or shared between atoms. The electrons that interact in this manner are those in the highest energy levels. These are the electrons most subject to the influence of nearby atoms or ions. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. Valence electrons are often located in incompletely filled main-energy levels. For example, the electron lost from the 3s sublevel of Na to form Na+ is a valence electron. For main-group elements, the valence electrons are the electrons in the outermost s and p sublevels. The inner electrons are in filled energy levels and are held too tightly by the nucleus to be involved in compound formation. The Group 1 and Group 2 elements have one and two valence electrons, respectively, as shown in Figure 3.10. The elements of Groups 13–18 have a number of valence electrons equal to the group number minus 10. In some cases, both the s and p sublevel valence electrons of the p-block elements are involved in compound formation (Figure 3.10). In other cases, only the electrons from the p sublevel are involved. Figure 3.10 Valence Electrons in Main-Group Elements 158 Chapter 5 Group number Group configuration Number of valence electrons 1 ns1 1 2 ns 2 2 13 ns 2p1 3 14 ns 2p 2 4 15 ns 2p 3 5 16 ns 2p 4 6 17 ns 2p 5 7 18 ns 2p 6 8 5C Main Idea Atoms have different abilities to capture electrons. Valence electrons hold atoms together in chemical compounds. In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another. This uneven concentration of charge has a significant effect on the chemical properties of a compound. It is therefore useful to have a measure of how strongly one atom attracts the electrons of another atom within a c ompound. Linus Pauling, one of America’s most famous chemists, devised a scale of numerical values reflecting the tendency of an atom to attract electrons. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. The most electronegative element, fluorine, is arbitrarily assigned an electronegativity of four. Other values are then calculated in relation to this v alue. Period Trends As shown in Figure 3.11, electronegativities tend to increase across each period, although there are exceptions. The alkali and alkaline-earth metals are the least electronegative elements. In compounds, their atoms have a low attraction for electrons. Nitrogen, oxygen, and the halogens are the most electronegative elements. Their atoms attract electrons strongly. Figure 3.11 Periodic Table of Electronegativities Shown are the electronegativities of the elements according to the Pauling scale. The most-electronegative elements are located in the upper right of the p-block. The least-electronegative elements are located in the lower left of the s-block. 1 Atomic number 1 H 2 3 5 6 7 Group 2 3 4 Li Be 11 12 1.0 C Group 13 Group 14 Symbol 2.5 1.5 2 5 B Electronegativity Na Mg 1.2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 19 20 21 22 23 24 25 26 27 0.9 K Ca Sc Ti 37 38 39 0.8 1.0 1.3 Group 10 Group 11 Group 12 28 29 30 6 C Group 15 Group 16 Group 17 7 N 8 O 9 Ne 18 2.5 3.0 3.5 4.0 13 14 15 16 17 Al Si 1.8 31 32 1.5 P — 10 F 2.0 — 2.1 S 2.5 Cl Ar 33 34 35 36 3.0 — 1.5 V 1.6 Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 3.0 Rb Sr 1.0 1.2 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te 2.1 2.5 I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 0.8 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.6 Cs Ba La Hf Ta W Re Os 2.2 2.2 Ir Pt Au Hg Tl Pb Bi Po At Rn 87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 0.7 Fr 0.7 0.9 Ra 0.9 1.1 Ac 1.1 1.3 Rf — 1.5 Db — 1.7 Sg — 1.9 Bh — Hs — Mt — 2.2 Ds — 2.4 Rg — 1.9 Cn 1.8 — 1.8 Fl 1.9 — 2.0 Lv 2.2 1 He 2.4 2 3 4 Period Period 4 6 2.1 Group 1 Group 18 5 6 7 — Lanthanide series 58 59 60 61 62 63 64 65 Ce Pr Nd Pm Sm Eu Gd Tb 90 91 92 93 94 95 96 97 1.3 1.5 Actinide series 1.4 1.1 Th 1.1 Pa 1.1 U 1.1 Np 1.4 1.2 Pu 1.3 1.1 Am 1.3 1.2 Cm 1.3 1.1 Bk 1.3 66 Dy 1.2 98 Cf 1.3 67 68 69 70 71 Ho Er Tm Yb Lu 99 100 101 102 103 1.3 — 1.2 Es 1.3 1.2 Fm 1.3 1.3 Md 1.3 1.1 No 1.3 Lr The Periodic Law 159 CHECK FOR UNDERSTANDING Electronegativities tend to either decrease down a group or remain about the same. The noble gases are unusual in that some of them do not form compounds and therefore cannot be assigned electronegativities. When a noble gas does form a compound, its electronegativity is rather high, similar to the values for the halogens. The combination of the period and group trends in electronegativity results in the highest values belonging to the elements in the upper right of the periodic table. The lowest values belong to the elements in the lower left of the table. These trends are shown graphically in Figure 3.12. Predict Metals generally have lower electronegativities than nonmetals. What type of ion would this suggest each would form? 5B Main Idea The properties of d -block metals do not vary much. The properties of the d-block elements (which are all metals) vary less and with less regularity than those of the main-group elements. This trend is indicated by the curves in Figures 3.3 and 3.5, which flatten where the d-block elements fall in the middle of Periods 4–6. Recall that atoms of the d-block elements contain from zero to two electrons in the s orbital of their highest occupied energy level and one to ten electrons in the d sublevel of the next-lower energy level. Figure 3.12 Electronegativity and Atomic Number The plot shows electronegativity versus atomic number for Periods 1–6. Electronegativity vs. Atomic Number F 4.0 Period 2 Period 3 Period 4 Period 5 Period 6 3.5 Kr Cl 3.0 Electronegativity Xe Rn 2.5 H 2.0 1.5 1.0 Li Na 0.5 0 10 K 20 Rb 30 Cs 40 Atomic number 160 Chapter 5 50 60 70 80 go online Periodic Trends in Electronegativity Solve It! Cards HMDScience.com Sample Problem G Of the elements gallium, Ga, bromine, Br, and calcium, Ca, which has the highest electronegativity? Explain your answer in terms of periodic trends. 5C SOLVE All of these elements are in the fourth period. Bromine has the highest atomic number and is farthest to the right in the period. Therefore, bromine should have the highest electronegativity because electronegativity increases across the p eriods. Answers in Appendix E 1. Consider five hypothetical main-group elements, E, G, J, L, and M, that have the outer electron configurations shown below. E = 2s22p5 G = 4d105s25p5 J = 2s22p2 L = 5d106s26p5 M = 2s22p4 a. Identify the block location for each element. Then, determine which elements are in the same period and which are in the same group. b. Which element would you expect to have the highest electron affinity? Which would you expect to form a 1- ion? Which should have the highest e lectronegativity? c. Compare the ionic radius of the typical ion formed by the element G with the radius of the atom from which the ion was formed. d. Which element(s) contains seven valence e lectrons? Therefore, electrons in both the ns sublevel and the (n–1)d sublevel are available to interact with their surroundings. As a result, electrons in the incompletely filled d sublevels are responsible for many characteristic properties of the d-block elements. Atomic Radii The atomic radii of the d-block elements generally decrease across the periods. However, this decrease is less than that for the main-group elements, because the electrons added to the (n–1)d sublevel shield the outer electrons from the nucleus. Also, note in Figure 3.3 that the radii dip to a low and then increase slightly across each of the four periods that contain d-block elements. As the number of electrons in the d sublevel increases, the radii increase because of repulsion among the e lectrons. In the sixth period, the f-block elements fall between lanthanum (Group 3) and hafnium (Group 4). Because of the increase in atomic number that occurs from lanthanum to hafnium, the atomic radius of hafnium is actually slightly less than that of zirconium, Zr, the element immediately above it. The radii of elements following hafnium in the sixth period vary with increasing atomic number in the usual manner. The Periodic Law 161 Ionization Energy As they do for the main-group elements, ionization energies of the d-block and f-block elements generally increase across the periods. In contrast to the decrease down the main groups, however, the first ionization energies of the d-block elements generally increase down each group. This is because the electrons available for ionization in the outer s sublevels are less shielded from the increasing nuclear charge by electrons in the incomplete (n–1)d sublevels. Ion Formation and Ionic Radii Among all atoms of the d-block and f-block elements, electrons in the highest occupied sublevel, the s and p sublevels, are always removed first. For the d-block elements, this means that although newly added electrons occupy the d sublevels, the first electrons to be removed are those in the outermost s sublevels. For example, iron, Fe, has the electron configuration [Ar]3d64s2. First, it loses two 4s electrons to form Fe2+ ([Ar]3d6). Fe2+ can then lose a 3d electron to form Fe3+ ([Ar]3d5). Most d-block elements commonly form 2+ ions in compounds. Some, such as iron and chromium, also commonly form 3+ ions. The Group 3 elements form only ions with a 3+ charge. Copper forms 1+ and 2+ ions, and silver usually forms only 1+ ions. As expected, the cations have smaller radii than the atoms do. Comparing 2+ ions across the periods shows a decrease in size that parallels the decrease in atomic radii. Electronegativity The d-block elements all have electronegativities between 1.1 and 2.5. Only the active metals of Groups 1 and 2 have lower electronegativities. The d-block elements also follow the general trend for electronegativity values to increase as radii decrease and vice versa. The f-block elements all have similar electronegativities, which range from 1.1 to 1.5. Section 3 Formative ASSESSMENT Reviewing Main Ideas 1.State the general period and group trends among main-group elements with respect to each of the following properties: 5C a.atomic radii b.first ionization energy c.electron affinity d.ionic radii e.electronegativity 2.a. In general, how do the periodic properties of the d-block elements compare with those of the main-group elements? 5B b.Explain the comparison made in ( a). 162 Chapter 5 3.For each main-group element, what is the relationship between its group number and the number of valence electrons that the group members have? Critical Thinking 4. RELATING IDEAS Graph the general trends (left to right and top to bottom) in the second ionization energy (IE2) of an element as a function of its atomic number, over the range Z = 3–20. Label the minima and maxima on the graph with the appropriate element symbol. 5C
