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UNIT IV – THE PERIODIC TABLE / PERIODICITY Periodicity Johann Dobereiner, 1817 • 1. Triad - group of three elements with similar properties. John Newlands, 1863 • • • 1. Arranged elements in order of their atomic masses. 2. Noticed that their properties repeated every 8th element. 3. Law of Octaves - the same properties repeated every 8th element Dmitri Mendeleev, 1869 • • Father of the Periodic Table Believed that similar properties occurred after periods that could vary in length: 1. First two rows: 7 elements 2. Second two rows: 17 elements Organized the periodic table according to the properties of the elements. Properties of the elements repeat in an orderly way. Such a repeating pattern is "periodic" Periodic Law - Properties of the elements are a periodic function of their atomic masses. Henry Moseley, 1911 • Worked with Ernest Rutherford and showed that the nucleus had a positive charge. He recognized the pattern of chemical properties when the elements were organized according to their nuclear charge (i.e., atomic number), Thus, the periodic law was revised: Periodic Law - Properties of the elements are a periodic function of their atomic numbers. The periodic table is the arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. Noble Gases were added to the periodic table later - early 1900's Lanthanides - f-Block was added in the early 1900's after the identity and properties of these elements (at. no. 58-71) were sorted out. Actinides - f-Block elements 90-103 were then added as they were identified Periodicity - can be observed by observing the properties of the elements either along any row (Period) or column (Group/Family). Electron Configuration and The Periodic Table Electron Configuration determines a chemical's reactivity. s-Block Elements (Groups 1 and 2; or Groups I A and II A) • • • Outermost electrons are added to an s-orbital Group 1: s1 - Alkali Metals Group 2: s2 - Alkaline Earth Metals p-Block Elements (Groups 13-18; Groups III A through VIII A) • • • • • • • Outermost electrons are added to a p-orbital Group 13: p1 Group 14: p2 Group 15: p3 Group 16: p4 - Chalcogens Group 17: p5 - Halogens Group 18: p6 - Noble Gases (Inert Gases) d-Block Elements (Groups 3-12; Groups I B through VIII B) • • • Outermost electrons are added to a d-orbital Known as the transition metals All metals Reading the electron configuration directly off the Periodic Table. • • • Main Group Elements - s-Block and p-Block Transition Elements - d-Block Lanthanides and Actinides - f-Block Elemental Properties: Metals • • • hard and shiny; conduct electricity on left side of the periodic table 1-3 electrons in outer level Nonmetals • • • gases or brittle solids; good insulators on right side of the periodic table >5 electrons in outer level Metalloids • properties of both metals and nonmetals Rule of Thumb have 1-3 electrons in Metals the outer shell have >5 electrons in Nonmetals the outer shell The Octet Rule - Eight electrons in the outer shell render an atom essentially unreactive. Rule of Thumb: An atom having a filled or half-filled sublevel is slightly more stable (less reactive) than an atom without a filed or half-filled sublevel Relative Atomic Stability (Decreasing order of stability) Full outer shell Full sublevel (s, p, d or f) half-filled sublevel no special arrangement of electrons If the last electron for an atom is in a full or half-full sublevel, then the atom is inherently more stable PERIODIC PROPERTIES We will look at the following periodic properties of the elements: 1 Atomic Radii 2 Ionic Radii 3 First Ionization Energy Oxidation Numbers / Valence 4 Electrons 5 Electronegativity 6 Electron Affinity 1. Atomic Radius As the principal quantum number (n) increases, the size of the electron cloud increases. That is, the atomic size increases as you go down the table. The reason for this is that you are adding energy levels as you go down the table (1, 2, 3,...). The positive charge of the nucleus increase as you go from left to right across the table. This increase in nuclear charge increase the pull on the electron cloud by the nucleus pulling the the electron cloud in tighter to the nucleus. Thus, the atom decreases in size. 2. Ionic Radius Metallic Ions • • Formed by the loss of electrons Smaller than the atoms from which they were formed Nonmetallic Ions • • Formed by the gain of electrons Larger than the atoms from which they were formed 3. First Ionization Energy • • • The energy required to remove the most loosely-held electron from a neutral atom. The first ionization energy tends to increase as atomic number increases in any horizontal row (or period). In any column or group, there is a gradual decrease in the first ionization energy as atomic number increases. Metals: low first ionization energies. Nonmetals: high first ionization energies Two factors which tend to lower the first ionization energy as you go down a particular group are: 1. Shielding effect- where inner electrons block the attraction of the nucleus for the outer electrons. 2. Increased distance of the outer electrons from the nucleus. The first ionization energy increases as you go across the periodic table. This is a result of an increasing nuclear charge Factors Affecting Ionization Energy • • • • Nuclear Charge - ionization energy is proportional to the nuclear charge Shielding Effect - Ionization energy is inversely proportional to the shielding effect Radius - Ionization energy is inversely proportional to the distance between the nucleus and the outer electrons Sublevel - an electron from a full or half-full sublevel requires additional energy to be removed 4. Oxidation Numbers Group 1 Group 2 Group 312 Group 13 Group 14 Group 15 Group 16 Group 17 Group 18 lose 1 electron lose 2 electrons multiple gain/loss lose 3 electrons lose/gain 4 electrons gain 3 electrons gain 2 electrons gain 1 electron stable +1 +2 +3 ±4 -3 -2 -1 0 5. Electronegativity Electronegativity - The relative tendency of an atom to attract electrons to itself when it is bonded to another atom Influenced by the same factors which affect ionization energy and electron affinity 1. Size 2. Shielding effect 3. Nuclear charge The Trends (in the Periodic Table) are the same • • increases from left to right increases from bottom to top The most active metals have the lowest electronegativity (Francium). The most active nonmetals have the highest electronegativity (Fluorine). Many chemical properties of the elements can be organized in terms of electronegativity. 6. Electron Affinity • • the attraction an isolated atom has for an additional electron shows the same trend as first ionization energy (increases from left to right, and decreases from the top down). Bond Strength - Two elements will react with each other depending on their relative attraction for electrons