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Ch. 5 - The Periodic Table C. Johannesson I II III Mendeleev Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were grouped together. Repeating patterns are referred to as periodic. Predicted properties of undiscovered elements. . Moseley Henry Mosely (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement. The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. The Modern Periodic Table • The Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. Visual Concept Ch. 5 - The Periodic Table II. Organization of the Elements C. Johannesson I II III Metallic Character Metals Nonmetals Metalloids Blocks Main Group Elements Transition Metals Inner Transition Metals Periodic Patterns Period of an element can be determined from elements electron configuration The name of each block (s,p,d &f) determined by what sublevel is being filled s 1 2 3 4 5 6 7 f (n-2) p d (n-1) 6 7 © 1998 by Harcourt Brace & Company C. Johannesson Periodic Patterns Period # energy level (subtract for d & f) A/B Group # total # of valence eColumn within sublevel block # of e- in sublevel Periodic Patterns Example - Hydrogen 1 2 3 4 5 6 7 1 1s 1st Period 1st column of s-block s-block Periods and Blocks of the Periodic Table, Sample Problem B An element has the electron configuration [Kr]4d55s2. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group. Ch. 5 - The Periodic Table Atomic Radius (pm) 250 III. Periodic Trends 200 150 100 50 0 0 5 10 Atomic Number 15 20 C. Johannesson I II III Trends Generally electron configuration of atom’s highest occupied energy level govern’s atom’s properties Vertical groups share similar chemical properties Horizontal rows/periods , length is determined by number of e- that can occupy sublevels being filled Atomic Radius Atomic Radius size of atom Increases to the LEFT and DOWN 1 2 3 4 5 6 7 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 positive nuclear charge without additional shielding pulls e- in tighter Ionization Energy Electron can be removed from an atom if enough energy is supplied Ion- an atom or group of bonded atoms that has a charge +/Ionization- process that results in formation of an ion Ionization energy (IE) energy required to remove one electron from a neutral atom Can compare the ease with which atoms give up electrons © 1998 LOGAL Ionization Energy First Ionization Energy Increases UP and to the RIGHT 1 2 3 4 5 6 7 Ionization Energy Ease of e- loss is major reason for high reactivity 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- Ionic Radius Ionic Radius Cations (+) lose esmaller Anions (–) gain e- larger C. Johannesson © 2002 Prentice-Hall, Inc. Ion Visual Concept Melting/Boiling Point Melting/Boiling Point Highest in the middle of a period. Increase as you go down 1 2 3 4 5 6 7 Examples Which atom has the larger radius? Beor Ba Ba Caor Br Ca Examples Which atom has the higher 1st I.E.? N or Bi N Baor Ne Ne Examples Which atom has the higher melting/boiling point? Li or C C Cr or Kr Cr Examples Which particle has the larger radius? S or 2S 2S Al or 3+ Al Al Ch. 5 - The Periodic Table Atomic Radius (pm) 250 III. Periodic Trends Continued 200 150 100 50 0 0 5 10 Atomic Number 15 20 C. Johannesson I II III Electron Affinity • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. • Electron affinity generally increases across periods. • Increasing nuclear charge along the same sublevel attracts electrons more strongly • Electron affinity generally decreases down groups. • The larger an atom’s electron cloud is, the farther away its outer electrons are from its nucleus. Electron Affinity Visual Concept Electronegativity Measure of the ability of an atom in chemical compound to attract eThere is an un even concentration of charge in a compound Effects chemical properties Fluorine most electronegative element Groups 1 & 2 least electronegative elements Tends to increase across periods decrease or remain the same down Electronegativity Visual Concept Example Of the elements gallium, Ga, bromine, Br, and calcium, Ca, which has the highest electronegativity? Explain your answer in terms of periodic trends. Bromine should have the highest electronegativity because electronegativity increases across the periods. Valence Electrons • Chemical compounds form because electrons are lost, gained, or shared between atoms. • The electrons that interact in this manner are those in the highest energy levels. • The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. • Valence electrons are often located in incompletely filled main-energy levels. • example: the electron lost from the 3s sublevel of Na to form Na+ is a valence electron. Valence Electrons Visual Concept
