Ch. 6 - The Periodic Table & Periodic Law I. Development of the Modern Periodic Table (p. 174 - 181) I II III A. Mendeleev Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass Elements with similar properties were grouped together There were some discrepancies A. Mendeleev Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered elements B. Moseley Henry Moseley (1913, British) Organized elements by increasing atomic number Resolved discrepancies in Mendeleev’s arrangement This is the way the periodic table is arranged today! C. Modern Periodic Table Group (Family) Period 1 2 3 4 5 6 7 1. Groups/Families Vertical columns of periodic table Numbered 1 to 18 from left to right Each group contains elements with similar chemical properties 2. Periods Horizontal rows of periodic table Periods are numbered top to bottom from 1 to 7 Elements in same period have similarities in energy levels, but not properties 3. Blocks Main Group Elements Transition Metals Inner Transition Metals 3. Blocks 1 2 3 4 5 6 7 Overall Configuration Lanthanides - part of period 6 Actinides - part of period 7 Ch. 6 - The Periodic Table II. Classification of the Elements (pages 182-186) I II III 1. Metals Good conductors of heat and electricity Found in Groups 1 & 2, middle of table in 3-12 and some on right side of table Have luster, are ductile and malleable a. Alkali Metals Group 1 1 Valence electron Very reactive Electron configuration ns1 Form 1+ ions Cations Examples: Li, Na, K b. Alkaline Earth Metals Group 2 Reactive (not as reactive as alkali metals) Electron Configuration ns2 Form 2+ ions Cations Examples: Be, Mg, Ca, etc c. Transition Metals Groups 3 - 12 Reactive (not as reactive as Groups 1 or 2), can be free elements Electron Configuration ns2(n-1)dx where x is column in d-block Form variable valence state ions Cations Examples: Co, Fe, Pt, etc 2. Nonmetals Not good conductors Found on right side of periodic table – AND hydrogen Usually brittle solids or gases a. Halogens Group 17 (7A) Very reactive Electron configuration ns2np5 Form 1- ions – 1 electron short of noble gas configuration Anions Examples: F, Cl, Br, etc b. Noble Gases Group 18 Unreactive, inert, “noble”, stable Electron configuration ns2np6 full energy level Have a 0 charge, no ions Examples: He, Ne, Ar, Kr, etc 3. Metalloids Sometimes called semiconductors Form the “stairstep” between metals and nonmetals Have properties of both metals and nonmetals Examples: B, Si, Sb, Te, As, Ge, Po, At C. Valence Electrons outermost s & p orbital electrons Stable octet - filled s & p orbitals (8 e-) in one energy level Group #A = # of valence e- (except He) 1A 1 2 3 4 5 6 7 8A 2A 3A 4A 5A 6A 7A C. Valence Electrons Valence electrons = electrons in outermost orbitals (highest principle energy level) You can use the Periodic Table to determine the number of valence electrons Each group has the same number of valence electrons 1A 1 2 3 4 5 8A 2A 3A 4A 5A 6A 7A Ch. 6 - The Periodic Table Atomic Radius (pm) 250 III. Periodic Trends (p. 187-194) 200 150 100 50 0 0 5 10 Atomic Number 15 20 I II III A. Periodic Law When elements are arranged in order of increasing atomic #, elements with similar chemical and physical properties appear at regular intervals. Atomic Radius (pm) 250 200 150 100 50 0 0 5 10 Atomic Number 15 20 B. Chemical Reactivity Families Similar valence e- within a group result in similar chemical properties 1 2 3 4 5 6 7 C. Other Properties Atomic Radius size of atom Ionization Energy © 1998 LOGAL Energy required to remove an e- from a neutral atom Electronegativity © 1998 LOGAL Shielding Effect There is a Nuclear charge experienced by the outer (valence) electron(s) in a multi-electron atom is due to the difference between the charge on the nucleus and the charge of the core electrons (inner electron shells). -Results in the reduction of attractive force between the positive nucleus and the outermost electrons due to “shielding effect” of the inner electron shells(core electrons). Periodic Trend, 1. Shielding effect increases down a group. 2. Shielding effect remains constant across a period. 1. Atomic Radius Atomic Radius = ½ the distance between two identical bonded atoms 1. Atomic Radius Atomic Radius Increases to the LEFT and DOWN 1 2 3 4 5 6 7 1. Atomic Radius Why larger going down? Higher energy levels have larger orbitals Shielding - core e- block the attraction between the nucleus and the valence e- Why smaller to the right? Increased nuclear charge without additional shielding pulls e- in tighter 2. Ionization Energy The minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. The ease with which an atom loses an e-. First Ionization Energy = Energy required to remove one e- from a neutral atom. He 1st Ionization Energy (kJ) 2500 Ne 2000 Ar 1500 1000 500 Na Li 0 0 5 10 Atomic Numbe r 15 K 20 2. Ionization Energy First Ionization Energy Increases UP and to the RIGHT 1 2 3 4 5 6 7 2. Ionization Energy Why opposite of atomic radius? In small atoms, e- are close to the nucleus where the attraction is stronger Why small jumps within each group? Stable e- configurations don’t want to lose e- 2. Ionization Energy Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed. Mg Core e- 1st I.E. 736 kJ 2nd I.E. 1,445 kJ 3rd I.E. 7,730 kJ 2. Ionization Energy Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed. Al Core e- 1st I.E. 577 kJ 2nd I.E. 1,815 kJ 3rd I.E. 2,740 kJ 4th I.E. 11,600 kJ Electron Affinity Most atoms can attract e- to form negatively charged ions The energy change that occurs when an e- is added to a gaseous atom or ion. The ease with which an atom gains an e-. For most atoms, the energy released when an e- is added. (in kJ/mol) Periodic Trend 1. Electron affinity slightly decreases down a group. 2. Electron affinity generally tends to increase across a period. 3. Electronegativity The measure of the ability of an atom in a chemical compound to attract electrons Given a value between 0 and 4, 4 being the highest Tendency for an atom to attract e- closer to itself when forming a chemical bond with another atom. 1 2 3 4 5 6 7 3. Electronegativity Why increase as you move right? More valence electrons, need less to fill outer shell Why increase as you move up? Smaller electron cloud, more pull by + nucleus Ionic Radius The size atoms become when losing or gaining electrons. Positive Ions – Metal - Atoms that lose e- and form positive ions become smaller. The lost e- is a valence e- and the atom may lose a shell.The repulsion between the remaining e- is lessened and allows the effective positive nuclear charge to pull the remaining ecloser. Negative Ions – Nonmetal - Atoms that gain e- and form negative ions become larger. The repulsion between the added e- and existing e- is increased and the effective positive nuclear charge cannot hold onto the e- tightly. Examples Which atom has the larger radius? Be or Ba Ca or Br Examples Which atom has the higher 1st I.E.? N or Bi Ba or Ne Examples Which element has the higher electronegativity? Cl or F Be or Ca More Practice Answer questions 16-19 on page 189 and 20-22 on page 194.