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CHAPTER 5, THE PERIODIC LAW Section 1, History of the Periodic Table By 1860, more than 60 elements had been discovered. FYI, there was no method for accurately determining an element’s atomic mass or the number of atoms of an element in a particular chemical compound. In 1869, Dimitri Mendeleev published his periodic table of elements in which elements with similar properties were grouped together. The elements were arranged in order of increasing atomic mass. He left empty spaces in the table. In 1871, he predicted the existence of elements that would fill these empty spaces. By 1886, all three elements had been discovered: Sc, Ga, and Ge. And, the properties were strikingly similar to those predicted by Mendeleev. Furthermore, Mendeleev is considered the discoverer of the periodic law. In the early 1900s, the chemistry of the lanthanides was understood. The lanthanides are the 14 elements with atomic number from 58 to 71. To follow was the discovery of the actinides. The actinides are the 14 elements with atomic number 90 to 103. Periodicity with respect to atomic number can be observed in any group of elements in the periodic table. Section 2, Electron Configuration and the Periodic Table Sublevel Blocks of the Periodic Table (image credit, modified: http://creationwiki.org/Periodic_table_blocks) KEY Mendeleev’s Periodic Table (image credit: http://www.chemistryexplained.com/NyPi/Periodic-Table.html) AM – alkali metal AEM – alkaliearth metal TM – transition metal M/S – metalloid/semimetal H – halogen NG – noble gas L – lanthanide A – actinide In 1911, Henry Moseley rearranged the periodic table of elements by putting elements in order of increasing nuclear charge, or number of protons in the nucleus. His work led to the modern definition of atomic number and the recognition that atomic number, not atomic mass, is the basis for the organization of the periodic table. Periodic law states that the physical and chemical properties of the elements are periodic function of their atomic numbers. 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. Between 1895 and 1898, Sir William Ramsay proposed adding the Group 18 elements, the noble gases, to the periodic table. Periodic Table of Elements (image credit, modified: http://intro.chem.okstate.edu/AP/2003Norman/HTML/SCH7.HTM) s-block • Group 1, alkali metals • Group 2, alkaline-earth metals • Hydrogen • Helium (noble gas) p-block • Groups 13-18 o Metalloids or Semiconductors o Group 17, halogens o Group 18, noble gases PERIODIC TRENDS Moving from left to right across a period Atomic Radii Decreases Why? Increasing positive charge of the nucleus. Ionization Energy Increases Why? Increasing nuclear charge. Electron Affinity Increases (becomes more negative) Why? Increasing nuclear charge allow for electrons to be more easily acquired. Ionic radii Decreases Why? The electron cloud shrinks due to increasing nuclear charge. Increases d-block • Groups 3-12, transition metals f-block • Lanthanides • Actinides Section 3, Electron Configuration and Periodic Properties Atomic radius is defined as ! the distance between the nuclei of identical atoms that are bonded together. An ion is an atom or group of bonded atoms that has a positive or negative charge. Ionization is any process that results in the formation of an ion. The atom’s electron affinity is the energy change that occurs when an electron is acquired by a neutral atom. A cation is a positive ion. The formation by loss of one or more electrons ALWAYS leads to a decrease in atomic radius because the removal of the highest-energy-level electrons results in a smaller electron cloud. And, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge. An anion is a negative ion. The formation by the addition of one or more electrons ALWAYS leads to an increase in atomic radius. This is because the total positive charge of the nucleus remains unchanged when an electron is added to an atom or an ion. So, the electrons are not drawn to the nucleus as strongly as they were before the addition of the extra electrons. The electron cloud also spreads out because of greater repulsion between the increased number of electrons. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Electronegativity Moving from top to bottom down a group Increases Why? Electrons occupy sublevels in successively higher main energy levels located farther away from the nucleus. Decreases Why? Additional energy levels and electron shielding. Electrons are farther from nucleus. Outer electrons are partially shielded from the effect of the nuclear charge by inner electrons. Decreases (becomes less negative) Why? Increasing atomic radius outweighs increasing nuclear charge, causing addition of electrons difficult. Increases Why? Occupationof successively higher energy levels. Decrease or remain about the same