I A 1 2 3 4 5 II A 1 2 3 4 Chapter 4 Periodic Trends of the Elements Development of the Periodic Table Historical Background John Newlands observed that when the elements were arranged according to increasing atomic mass every eight element exhibited similar properties He called this the law of octives Dmitri Mendeleev arranged elements according to their properties and was able to predict the properties of yet undiscovered elements gallium Henry Moseley a b mass Determined the atomic number of elements and found that with a few exceptions atomic number increases in the same order as atomic Concluded that atomic number was equal to the number of protons in the nucleus of an atom In the modern periodic table elements are arranged according to increasing atomic number and elements with similar properties appear in the same vertical column Name two elements whose properties are similar to sodium The Modern Periodic Table Classification of Elements Main Group or representative elements are the Groups 1A 7A which are characterized by incompletely filled s or p orbitals Transition metals are Groups 1B and 3B through 8B d block These elements are characterized by incompletely filled d orbitals Group 2B elements are neither representative nor transition elements because their d subshells are completely filled Lanthanides and actinides also known as the f block elements have incompletely filled f orbitals 1 B 1 2 3 C A 1 2 B 1 2 3 4 5 6 General Information about the Periodic Table Groups are the vertical columns in the periodic table Elements in a group have similar physical and chemical properties Periods are the horizontal rows in the periodic table Across a row electrons are added to the orbitals of a specific principal quantum level and end with a noble gas Valence electrons are the outer most electrons in an atom These are the electrons that are involved in chemical bonding Electron Configuration and Group Number Group 1A Li He 2s1 Na Ne 3s1 K Ar 4s1 Rb Kr 5s1 Cs Xe 6s1 Group 2A Be He 2s2 Mg Ne 3s2 Ca Ar 4s2 Sr Kr 5s2 Ba Xe 6s2 Group 5A N He 2s2 2p3 P Ne 3s2 3p3 As Ar 3d10 4s2 4p3 Sb Kr 4d10 5s2 5p3 Bi Xe 4f145d10 6s26p3 Note that the group number corresponds to the number of valence electrons III Effective Nuclear Charge Definitions Nuclear charge Z is determined by the number of protons in the nucleus Effective nuclear charge Zeff is the magnitude of positive charge felt by an electron in an atom Details of effective nuclear charge The only element in which Z and Zeff are the same is hydrogen In all other atoms the electrons are simultaneously attracted by the nucleus and repelled by the other electrons This phenomenon is known as shielding In a multi electron atom electrons are shielded from the positive charge of the nucleus by other electrons Core electrons are the electrons in completely filled inner shells These are the most effective at shielding Zeff increases from left to right across a period The number of core electrons remain the same while the number of protons in the nucleus increase The change in Zeff is less significant within a group than in a period 2 A 1 2 3 4 5 B 1 Zeff Z where is the shielding constant IV Periodic Trends in Properties of the Elements Atomic Radius Atomic radius can be defined as one half the distance between the nuclei of two adjacent identical metal atoms or one half the distance between the nuclei of two adjacent identical atoms joined by a chemical bond Atomic radius decreases from left to right across a period because the effective nuclear charge increases and the valence shell is drawn closer to the nucleus As protons are added to the nucleus and electron to the valence shell in each successive element there is a stronger attraction between the nucleus and the valence shell due to the higher nuclear charge As the nuclear charge increases atomic radius decreases Atomic radius increases from top to bottom within a group because electrons are filling successively higher principal quantum levels and are at a larger distance from the nucleus Referring only to the periodic table arrange the elements Ge Se and F in order of increasing atomic radius On the basis of their position in the periodic table select the atom in each pair that has the larger atomic radius a Mg P b Sr Be c As Br d Cl I e Xe Kr Ionization Energy IE Ionization energy is the minimum energy required to remove an electron from a gaseous atom in its ground state Ionization energy is expressed in kJ mol Na g Na g e 495 8 kJ mol 3 2 3 4 5 6 a b c d e This is the first ionization energy IE1 Ionization energy increases from left to right across a period and decreases from top to bottom within a group When an electron is removed from a neutral atom a positive ion forms that is known as a cation First ionization energy as a function of atomic number Note that Group 1A elements have the lowest and the noble gases have the highest ionization energies It is possible to remove a second or a third electron but each successive ionization requires more energy IE1 IE2 IE3 Once an ion has a noble gas configuration there is a dramatic increase in the amount of energy required to remove a core electron Refer to Table 4 2 Practice Exercises Circle the atom in each pair that has the smaller first ionization IE1 O or S K or Ca Which atom has larger IE2 Li or Be Which atom has larger IE1 N or P Which atom has smaller IE2 Na or Mg At which IE should a dramatic increase in energy occur for Al 4 C Electron Affinity EA 1 2 3 4 5 6 7 Electron affinity EA is the energy change that occurs when an electron is added to a gaseous atom For most atoms energy is released when an electron is added Electron affinity is expressed in kJ mol and the negative sign indicates that energy is released F g e F g EA 328 kJ mol When 1 mole of F atoms accept one mole of electrons 328 kJ of energy is released When an electron is added to a neutral gaseous atom a negative ion forms that is known as an anion Generally EA becomes more negative from left to right across a period due to the increase in Zeff There are deviations from these trends that are due to the electron configuration of the atoms to which electron is added For example it is more difficult to add an electron to a Group 2A element than to a Group 1A because the electron must go into a higher energy orbital in 2A See figure below More than one electron can be added to an atom Many first EA s are negative while subsequent EA s are positive The input of energy is