<p> 64</p><p>Chapter 8: Periodic Properties of the Elements</p><p>I) Preliminary Questions</p><p>A) What is periodicity? Why is it so important to chemistry? B) How can periodicity be used to explain observed trends in properties of groups of elements?</p><p>II) History of Periodic Table</p><p>A) J. Dobreiner</p><p>- Law of Triads - groups of three elements share similar properties.</p><p>B) J. Newlands</p><p>- Law of Octaves - elements arranged in octaves (did not account for Group VIIIA). - works well for s & p-block elements only.</p><p>C) Dmitri Mendeleev & J. Lothar Meyer</p><p>- Meyer based his periodic table on physical properties of elements.</p><p>- Mendeleev based periodic table on chemical properties of elements and went so far as to make predictions concerning the properties of undiscovered elements. (see Figure 8.1)</p><p>- Both based their periodic tables on increasing atomic mass.</p><p>- Modern table based on increasing atomic number.</p><p>3) Concept of Periodicity </p><p>A) Properties of certain elements are similar & occur at fixed intervals.</p><p>B) Periodicity used to predict trends due to electron configurations. (Example: Alkali Metals) 65</p><p>III) Electronic Configurations</p><p>1) Aufbau Principle (Diagonal Rule)</p><p>A) accounts for how electrons are introduced into atomic orbitals.</p><p>1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p</p><p>2) Hund’s Rule</p><p>A) electrons fill degenerate orbitals halfway first (results in lowest energy configuration).</p><p>B) electrons listed as having parallel spins.</p><p>3) Orbital Filling Diagrams </p><p>A) An alternate method of expressing electron configurations, where arrows represent the electrons. </p><p>B) Within a given orbital, electrons are paired with opposing spins (one arrow points up, the other points down).</p><p>4) Valence Electrons</p><p>A) electrons found in outermost principal quantum level of atom. Other electrons are called core electrons.</p><p>B) electrons in same family (group) of periodic table have same outer electronic configuration and share similar chemical behavior.</p><p>Examples: Group IA ns1</p><p>Group IIA ns2</p><p>Group IIIA ns2np1 66</p><p>C) 4 Blocks of Elements in Periodic Table</p><p>1. s-block main group (representative) elements</p><p>2. p-block main group (representative) elements</p><p>3. d-block transition metals</p><p>4. f-block inner transition metals</p><p>Table 1. Distribution of Electrons Inside Electron Cloud</p><p>Principal Energy Level (n) 1 2 3 4 Maximum No. of Electrons (2n2) 2 8 18 32 Atomic Orbitals in Sublevel 1s 2s, 2p 3s, 3p, 3d 4s,4p,4d,4f Designation of Filled Orbitals 1s2 2s2, 2p6 3s2, 3p6, 3d10 4s2, 4p6, 4d10, 4f14 Maximum Electrons per Sublevel 2 2, 6 2, 6, 10 2, 6, 10, 14 Orbitals per Sublevel 1 1, 3 1, 3, 5 1, 3, 5, 7</p><p>5) Pauli Exclusion Principle</p><p>A) no two electrons in an atom can have the same set of 4 quantum numbers.</p><p>6) Polyelectronic Atoms</p><p>A) atoms with more than 1 electron. B) solutions to Schroedinger equation are much more complicated.</p><p>IV) Electronic Configurations & Periodic Properties: Atomic Radii</p><p>1) Atomic Radius (Size)</p><p>- atomic radius decreases across period. - atomic radius increases down a column.</p><p>2) Ionization Energy</p><p>General Equation: X (g)  X + (g) + e -</p><p>A) Ionization energy represents the energy required to remove an electron from a gaseous atom (ion).</p><p>B) Ionization energy increases across period.</p><p>C) Ionization energy decreases down a column. 67</p><p>3) Electron Affinity</p><p>General Equation X (g) + e -  X - (g)</p><p>A) Electron affinity refers to the energy change associated with addition of an electron to a gaseous atom.</p><p>B) Electron affinity increases across period.</p><p>C) Electron affinity decreases down a column.</p><p>V) Introduction to Chemical Bonding</p><p>1) The chemical forces that hold atoms together are referred to as bonds. 2) Two main categories of bonding</p><p>A) Ionic Bonding B) Covalent Bonding</p><p>VI) Ions & Electronic Configurations</p><p>1) Valence shell electrons are what participate in chemical bonding. 2) Metals lose electrons to form cations, while nonmetals gain electrons to form anions. 3) Atoms want to attain a noble gas electron configuration. Since the noble gases (except He) have a total of 8 valence electrons, this condition is often called the octet rule. 4) For main group elements, the group # gives the number of valence electrons directly. (Group IA has 1 valence electron, Group IIA has 2 valence electrons, etc.). 5) Predicting Formulas for Ionic Compounds </p><p>Ionic Compounds  dealing with solids</p><p>Example: Calcium oxide CaO</p><p>Ca [Ar] 4s2 Ca2+ [Ar] O [He] 2s2p4 O2- [Ne]</p><p>Mention Exceptions that are found along diagonal (metalloids)</p><p>Isoelectronic ions: ions with same number of electrons.</p><p>Size decreases with increasing Z for a set of isoelectronic ions. 68</p><p>VII) Ionic Bonds & Formation of Ionic Solids</p><p>1) Ionic bonding involves a transfer of electrons from a metal atom to a nonmetal atom. This transfer results in the formation of metal cations and nonmetal anions.</p><p>2) All salts are ionic compounds. These compounds consist of a three dimensional solid consisting of alternating cations and anions.</p><p>3) We will not concern ourselves with Born-Haber Cycle or lattice energy calculations.</p>