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The Periodic Table Elements are arranged in a way that shows a repeating, or periodic, pattern. Dmitri Mendeleev created the first periodic table of the elements in 1869. He ordered the ~70 known elements by their atomic masses and their chemical properties. He found that some elements could not be put into groups with similar properties and at the same time stay in order. Mullis 1 Modern Periodic Table Later, Henry Moseley carried on the work. Moseley put the elements in order of increasing atomic NUMBER. He found that the position of the element corresponded to its properties. The modern periodic table shows the position of the element is related to : Atomic number AND Arrangement of electrons in its energy levels Mullis 2 Electron Shells Move down P. table: Principal quantum number (n) increases. Distribution of electrons in an atom is represented with a radial electron density graph. Radial electron density is probability of finding an electron at a particular distance from the nucleus. Electron shells are diffuse and overlap a great deal. Mullis 3 Examples of Electron shells He: 1s2 Radial plot shows 1 maximum Ne: 1s2 2s2 2p6 Radial plot shows 2 maxima ( 1 each for the 1st and 2nd energy levels ) Ar: 1s2 2s2 2p6 3s2 3p6 Radial plot shows 3 maxima ( 1 each for the 1st,2nd and 3rd energy levels ) Mullis 4 Atomic Sizes: Single atoms Colliding argon atoms ricochet apart because electron clouds cannot penetrate one another to a significant extent. The apparent radii are determined by the closest distances separating the nuclei during such collisions. This radius is called the nonbonding radius. Mullis 5 Atomic Sizes: Bonded atoms The distance between two nuclei is called the bond distance. If the two atoms making up the molecule are the same, then ½ the bond distance is called the bonding atomic radius of the atom. This radius is shorter than the nonbonding radius. Mullis 6 Atomic Sizes using Periodic Table As we move down a group, atoms become larger. Larger n = more shells = larger radius As we move across a period, atoms become smaller. More protons = more effective nuclear charge, Zeff More positive charge increases the attraction of nucleus to the electrons in the outermost shell, so the electrons are pulled in more “tightly,” resulting in smaller radius Mullis 7 Ionization energy Ionization energy of an ion or atom is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. The first ionization energy, I1 is the energy required to remove one electron from an atom. Na(g) Na+(g) + e The 2nd ionization energy, I2, is the energy required to remove an electron from an ion. Na+(g) Na2+(g) + e Larger ionization energy, harder to remove electron. Mullis 8 Periodic Trends in Ionization Energy Highest = Fluorine Ionization energy decreases down a group. Easier to remove electrons that are farther from the nucleus. Ionization energy increases across a period. Zeff increases, so it’s harder to remove an electron. Exceptions: Removing the 1st and 4th p electrons Mullis 9 Electron Affinity Gain Lose Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion. Electron affinity: Cl(g) + e- Cl-(g) Ionization energy: Cl(g) Cl+(g) + eAffinity for reaction above is exothermic: ∆E = -349 kJ/mol If adding the electron makes the species more stable, it will be exothermic. Mullis 10 Coulomb’s law Which law can best be used to explain why addition of an electron to the O2– ion is an endothermic process? Coulomb’s law: The energy required for the process is necessary to overcome the electrostatic repulsion between the electron and the already negatively charged O2– ion. Mullis 11 Ion size The oxide ion is isoelectronic (has exactly the same number and configuration of electrons) with neon, and yet O2– is bigger than Ne. Why? This is Coulomb's law at work. In any isoelectronic series the species with the highest nuclear charge will have the smallest radius. Mullis 12 Metals Metallic character increases down a group and from left to right across a period. Metal properties: Lustrous (shiny) Malleable (can be shaped) Ductile (can be pulled into wire) Conduct electricity Metal oxides form basic ionic solids: Metal oxide + water metal hydroxide Metal oxides react with acids to form salt and water Mullis 13 Metals Metal oxides form basic ionic solids: Metal oxide + water metal hydroxide MgO(s) + H2O(l) Mg(OH)2(s) Metal oxides react with acids to form salt and water MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l) Most neutral metals are oxidized rather than reduced. Metals have low ionization energies. Mullis 14 Metal reactivity Which of the alkali metals would you expect to react most violently with water? Li, Na, K, Rb Of these four, rubidium has the lowest ionization energy, making it the most reactive. Rubidium reacts explosively with water. Mullis 15 Nonmetals Lower melting points than metals Diatomic molecules are nonmetals. Most nonmetal oxides are acidic: Nonmetal oxide + water acid P4O10(s) + 6H2O(l) 4H3PO4(aq) Nonmetal oxides react with bases to form salt and water: CO2(g) + 2NaOH(aq) Na2CO3(aq) + H2O(l) Mullis 16 Nonmetallic oxides Which nonmetallic oxide would you expect to be the strongest acid? NO2, N2O, N2O4, N2O5 N2O5: Nitrogen has an oxidation state of +5 in this compound. In general, the higher the oxidation state of the nonmetal, the more acidic the nonmetal oxide. Mullis 17