Chapter 2.4 Periodic properties of the elements Electron Configurations, Atomic Properties, and Periodic Trends As we move down the periodic table the electron configuration becomes longer and more complex. This goes beyond the purpose of our course. Electron configurations help to determine the atomic and chemical properties of the elements. Properties such as atomic radius, ionization energy, electron affinity and electronegativity are periodic as they follow recurring trends in the periodic table. increase the orbital energy in oxygen. Therefore, less energy is required to remove the fourth p sublevel electron in oxygen. Periodic trends in ionization energy are linked to trends involving the reactivity of metals. In general, the chemical reactivity of metals increases down a group and decreases across a period. These trends, as well as a further trend from metallic to non-metallic properties across a period, and increasing metallic properties down a group, are shown in Table 3.1. Table 3.1 Periodic Trends in Reactivity of Metals With oxygen Reaction with Reaction in the air With water dilute acids Group 1 React rapidly to form React vigorously The reaction is oxides with the with cold water, dangerously general formula M2O. generating H2(g) violent, igniting These compounds and forming a H2(g) generated. are strong bases. strong base with the general formula MOH. Group 2 React moderately to React slowly in React quickly form basic oxides cold water, generating H2(g). with the general generating H2(g) Electron Configurations,formula MO. Atomicand forming a Properties, and strong base, Periodic Trends M(OH)2. Group 13 React slowly. Al displaces H2(g) Al reacts readily PeriodicAluminum Trends forms an in Reactivityfrom steam. of Metalsif the protective Reaction Withoxide oxygen coating in thatthe air Withoxide water coating is Reaction with dilute protects the metal; removed, to acids Group 1 React rapidlythe to compoundform oxides Al with2O3 the React vigorouslygenerate with cold H 2(g)water, . The reaction is dangerously can act as a base or general formula M2O. These compounds generating H2(g) and forming a violent, igniting H2(g) are strong bases.an acid. strong base with the general generated. Group 14 C and Si are non- The mostformula metallic MOH. Sn and Pb react Group 2 React moderatelymetals, to and form form basic oxideselements, React Sn slowly and in coldslowly water, to form React quickly generating with the generalacidic formula oxides. MO. Ge Pb, dogenerating not react H2(g)H and2(g) .forming a H2(g). is a metalloid, and with water. strong base, M(OH)2. forms an acidic oxide. Group 13 React slowly. Aluminum forms an oxide Al displaces H2(g) from steam. Al reacts readily if the SnO can act as an coating that protects2 the metal; the protective oxide coating is acid or a base. PbO2 compound Alis2 Ounreactive.3 can act as a base or an removed, to generate H2(g) . acid. Group 14 TransitionC and Metals Si areZn non- is themetals, most and form acidicZn and The Fe most metallicZn elements, reacts readily Sn Sn and Pb react slowly to (for example, Fe, reactive transition displace H2(g) to generate H2(g), oxides. Ge is a metalloid, and forms an and Pb, do not react with water. form H2(g). Co, Ni, Cu, Zn, metal, and forms ZnO from steam. Fe Cu, Ag, and Au acidic oxide. SnO2 can act as an acid or a Ag, Au) when the metal is rusts slowly at do not react. base. PbO is unreactive. 2 burned in air. Ag and room temperature. Transition Metals (for Zn is the mostAu reactivedo not react. transition metal, Zn and Fe displace H2(g) Zn reacts readily to generate example, Fe, Co, Ni, Cu, and forms ZnO when the metal is burned from steam. Fe rusts slowly at H2(g), Cu, Ag, and Au do not Zn, Ag, Au) in air. Ag and Au do not react. room temperature. react. Magnesium ribbon reacting Potassium reacting Nickel reacting with with oxygen in air with water dilute acid Chapter 3 Atoms, Electrons, and Periodic Trends • MHR 155 Periodic Trends in Atomic Radius As you know, atoms are not solid spheres. Their volumes (the extent of their orbitals) are described in terms of probabilities. Nevertheless, the size of an atom — its atomic radius — is a measurable property. Chemists can determine it by measuring the distance between the nuclei of bonded, neighbouring atoms. For example, for metals, atomic radius is half the Atomic Radius distance between neighbouring nuclei in a crystal of the metal element. For elements that occur as molecules, which is the case for many non- metals, atomic radius is half the distance between nuclei of identical atoms that bonded together with a single covalent bond. In Figure 3.23, the radii of metallic elements represent the radius of an atom in a metallic The size of an atom, its atomic radius, decreases crystal.across The radiia period. of all other elementsFurthermore, represent the radius atomic of an atom radii of the element participating in a single covalent bond with one additional, generally increase down a group. Two factors affectlike atom. differences in atomic radii: 1 8 (IA) (VIIIA) H 37 He 31 1 2 3 4 5 6 7 (IIA) (IIIA) (IVA) (VA) (VIA) (VIIA) Li152 Be 112 BC85 77 N 75 OF73 72 Ne 71 2 1.As n increases, there is a higher probability of Na186 Mg 160 Al143 Si 118 P 110 S103 Cl 100 Ar 98 finding electrons farther from their nucleus. 3 Therefore, the atomic volume is larger. K227 Ca 197 Ga135 Ge 122 As 120 Se119 Br 114 Kr 112 4 Period 2.The other factor that affects atomic radii is Rb248 Sr 215 In167 Sn 140 Sb 140 TeI142 133 Xe 131 changing nuclear charge — specifically, the 5 effective nuclear charge (the net force of Cs265 Ba 222 Tl170 Pb 146 Bi 150 Po 168At (140) Rn(140) attraction between electrons and the nucleus) 6 Fr (270)Ra (220) 7 3 4 5 6 7 8 9 10 11 12 (IIIB) (IVB) (VB) (VIB) (VIIB) (VIIIB) (IB) (IIB) Sc162 Ti 147 V134 Cr 128 Mn127 Fe 126 Co 125 Ni124 Cu 128 Zn 134 4 Y180 Zr 160 Nb146 Mo 139 Tc136 Ru 134 Rh 134 Pd137 Ag 144 Cd 151 5 La 187 Hf 159 TaW146 139 Re137 Os 135 Ir 136 Pt138 Au 144 Hg 151 6 Figure 3.23 Representations of atomic radii for main group and transition elements. (Values for atomic radii are given in picometres. Those in parentheses have only two significant digits.) 152 MHR • Unit 2 Structure and Properties First Ionization Energy The energy needed to completely remove one electron from a ground state gaseous atom is called the ionization energy. Energy must be added to the atom to remove an electron in order to overcome the force of attraction exerted on the electron by the nucleus. Since multi-electron atoms have two or more electrons, they also have more than one ionization energy. For calcium the first ionization energy (IE1), is 599 kJ/mol: Ca(g) + 599 kJ → Ca+(g) + e- The second ionization energy (IE2) is the amount of energy required to remove the second electron. For calcium, it may be represented as: Ca+(g) + 1145 kJ → Ca2+1(g) + e- For a given element, IE2 is always greater than IE1 because it is always more difficult to remove a negatively charged electron from a positively charged ion than from the corresponding neutral atom. Ionization energies measure how strongly electrons are bound to atoms. Ionization always requires energy to remove an electron from the attractive force of the nucleus. Low ionization energies indicate easy removal of electrons, and hence easy positive ion (cation) formation. First Ionization Energy Every element exhibits a large increase in ionization energy when an inner electron is removed. This supports the idea that only the outermost electrons are involvedSECTION 7.4in theIonization chemical Energy 261 bondings andGO reactions. FIGURE The inner electrons are too tightly bound to the nucleusWhich to hasbe a lostlarger firstfrom ionization the energy, atom Ar or orAs? Why?even shared with another atom. H 1312 Be 899 Li He 520 2372 Mg Na 738 496 Ne Ca K Sc 2081 419 590 Ti N F 633 1402 1681 659 V C O Rb Sr 651 Cr B 1086 403 549 Y Mn 1314 600 Zr 653 717 Fe 801 640 Nb 763 Co Cs Ba 652 Mo Lu 684 Tc 760 Ni P Ar Ionization energy (kJ/mol) 376 503 Hf Ru Si Cl 1521 524 659 Ta 702 737 Cu Zn Al 1012 S 761 W 710 Rh 746 906 578 786 1251 770 720 Pd 1000 Re Os 760 804 Ag Cd Ge As Kr 840 Ir 731 Ga 947 Se Br 1A 880 Pt 868 579 762 1140 1351 2A 870 Au Hg 941 890 1007 In Sn Sb 558 709 834 Te I Xe 869 1008 1170 Tl Pb 589 716 Bi Po 703 Rn 812 1037 Increasing ionization energy 3A 4A 5A 6A 7A Increasing ionization energy 8A į FIGURE 7.9 Trends inTrends first ionization in first energies ionization of the elements. energies of the elements. between the electron and the nucleus. As this attraction increases, it becomes more difficult to remove the electron and, thus, the ionization energy increases. As we move across a pe- riod, there is both an increase in effective nuclear charge and a decrease in atomic radius, GO FIGURE causing the ionization energy to increase. As we move down a column, the atomic radius in- Explain why it is easier to remove creases while the effective nuclear charge increases rather gradually.