Ch 5 Chemical Periodicity Ch 5 end of book chapter HW Q 4 7 10 11 13 16 22 24 25 28 30 32 36 3842 46 49 54 55 66 69 73 74 86 1 More About the Periodic Table Periodic Properties of the Elements 2 Atomic Radii 3 Ionization Energy 4 Electron Affinity 5 Ionic Radii 6 Electronegativity 7 Oxidation States Chemical Reactions and Periodicity 8 Hydrogen the Hydrides 9 Oxygen the Oxides 2 Ch 5 Chemical Periodicity More About the Periodic Table The properties of the elements are periodic functions of their atomic Establish a classification scheme of the elements based on their electron configurations 3 Noble Gases All of them have completely filled electron shells n2 np6 Since they have similar electronic structures their chemical reactions are similar He 1s2 Ne He 2s2 2p6 Ar Ne 3s2 3p6 Kr Ar 4s2 3d10 4p6 Xe Kr 5s2 4d10 5p6 Rn Xe 6s2 4f14 5d10 6p6 Representative Elements Are the elements in A groups on periodic chart These elements will have their last electron in an outer s or p orbital These elements have fairly regular variations in their properties 5 d Transition Elements Elements on periodic chart in B groups Sometimes called transition metals Each metal has d electrons ns n 1 d configurations These elements make the transition from metals to nonmetals Exhibit smaller variations from row to row than the representative elements 6 f transition metals Sometimes called inner transition metals Electrons are being added to f orbitals Electrons are being added two shells below the valence shell Consequently very slight st 1 variations of properties from f transition series lanthanides 4 f orbitals one element to another Outermost electrons have 58Ce thru 71Lu the greatest influence on the 2nd series actinides 5 f chemical properties of orbitals 90Th 103Lr elements 7 The outermost e have the greatest influence on the properties of elements Adding an e to an s or p orbital usually causes dramatic changes in the physical and chemical properties Adding an e to a d or f orbital typically has a much smaller effect on properties Outermost e are those that have the highest value of the principal quantum number n Periodic Properties of the Elements Atomic Radii Ionization Energy Electron Affinity Ionic Radii Electronegativity 9 Atomic Radii Atomic radii describes the relative sizes of atoms Atomic radii increase within a column going from the top to the bottom of the periodic table Atomic radii decrease within a row going from left to right on the periodic table This last fact seems contrary to intuition How does nature make the elements smaller even though the electron number is increasing 10 11 Atomic Radii The reason the atomic radii decrease across a period is due to shielding or screening effect Effective nuclear charge Zeff experienced by an electron is less than the actual nuclear charge Z The inner electrons block the nuclear charge s effect on the outer electrons Moving across a period each element has an increased nuclear charge and the electrons are going into the same shell 2s and 2p or 3s and 3p etc Consequently the outer electrons feel a stronger effective nuclear charge For Li Zeff 1 For Be Zeff 2 12 Atomic radii increase going down a group because e are being added to shells farther from the nucleus Atomic radii decrease from left to right within a given period owing to increasing effective nuclear charge H atoms are the smallest and Cs atoms are largest naturally occurring atoms Electrons in inner shells screen or shield e in outer shells from the full effect of the nuclear charge This concept of a screening or shielding effect helps us understand many periodic trends in atomic properties So each outer shell e is shielded by just 2 inner e B Z 5 1s2 2s2 2p1 C Z 6 1s2 2s2 2p2 N Z 7 1s2 2s2 2p3 O Z 8 1s2 2s2 2p4 F Z 9 1s2 2s2 2p5 Ne Z 10 1s2 2s2 2p6 Atomic Radii Arrange these elements based on their atomic radii Se S O Te O S Se Te Arrange these elements based on their atomic radii Ga F S As F S As Ge Clicker Q Arrange these elements based on their atomic radii P Cl S Si Cl S P Si 15 Ionization Energy First ionization energy IE1 The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1 ion Measures how tightly e are bound to atoms Low ionization energies indicate easy removal of e and hence easy positive ion formation Symbolically Atom g energy ion g e 16 For a given element IE2 is always greater than IE1 because it is always more difficult to remove a neg charged e from a pos charged ion than from the corresponding neutral atom The 1st e added to a shell is easily removed to form a noble gas configuration as for alkali metals Ionization Energy So as atomic radii increase in a given group IE1 decrease because the outermost e are farther from the nucleus 1st IE1 generally increase from left to rt across the periodic table because the e are held more tightly by nucleus Mg g 738kJ mol Mg e 18 Ionization Energy Second ionization energy IE2 The amount of energy required to remove the second electron from a gaseous 1 ion Symbolically ion energy ion2 eMg 1451 kJ mol Mg2 e 19 Ionization Energy Periodic trends for Ionization Energy 1 IE2 IE1 It always takes more energy to remove a second electron from an ion than from a neutral atom 2 IE1 generally increases moving from IA elements to VIIIA elements Important exceptions at Be Mg N P etc due to filled and half filled subshells 3 IE1 generally decreases moving down a family IE1 for Li IE1 for Na etc 20 First Ionization Energies of Some Elements 21 Ionization Energy Arrange these elements based on their first ionization energies Sr Be Ca Mg Be Mg Ca Sr Arrange these elements based on their first ionization energies B O Be N B Be O N 22 Clicker Q Arrange these elements based on their first ionization energies Al Cl Na P 23 Ionization Energy First second third etc ionization energies exhibit periodicity as well Look at the following table of ionization energies versus third row elements Notice that the energy increases enormously when an electron is removed from a completed electron shell 24 Ionization Energy Group and element IA Na IIA Mg IIIA Al IE1 kJ mol IE2 kJ mol IE3 kJ mol IE4 kJ mol 496 738 578 4562 1451 1817 6912 7733 2745 9540 10 550 11 580 25 Ionization Energy The reason Na forms Na and not Na2 is that the energy difference between IE1 and IE2 is so large Requires more than 9 times more energy to remove the second electron than the first one The same trend is persistent