The Known Elements. Periodic Trends: Valence Electrons. History of the. Trends on the. Periodic Table

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1 Trends on the Periodic Table History of the Periodic Table Mendeleeve (1869) : Ordered elements according to increasing molecular weight. Same groups had similar properties predicted presence of unknown elements The Known Elements Moseley ( ): Discovered atomic number concept and noticed that Mendeleeve s Periodic Table increased in atomic number. He found greater order existed and that the elements would fall in more cohesive groups. Mendeleeve could predict the missing weights, but Moseley found by using atomic number, elements could be more ordered Periodic Law When elements are arranged by increasing atomic numbers, groups display similar characteristics between elements. Valence Electrons Atoms in the same group have the same ground state electron configuration; increasing only in principle quantum number. Atoms in the same group have the same number of valence electrons.

2 1st Ionization Energy (kj/mol) 2nd Ionization Energy (kj/mol) (pm) Is taken as the covalent radius for non-metallic elements and as the metallic radius for metals Covalent radius = One-half the distance between the nuclei of two identical atoms that are singly bonded to one another. For covalent compounds whose atoms do not bond to one another, comparisons between atoms in multiple molecules is necessary Metallic radius is one-half the closest internuclear distance in a metallic crystal. Trends Increases down group and decreases across period Ionization Energy Is the minimum energy required to remove the outermost electron from an isolated gaseous atom or ion in the ground state. First Ionization Energy Energy required to remove the first electron from a neutral atom Second Ionization Energy Energy needed to remove the outermost electron from a +1 ion. Energy needed to remove the second electron from a neutral atom

3 Electron Affinity (kj/mol) Ionization Energy Trends Ionization energy decreases down the group and increases across the period. Also, each successive ionization energy is higher. Representative elements differ in larger magnitudes than transition metals. Electron Affinity Energy released when an electron is added to the valence level of a gas-phase atom. For most atoms energy is released when an electron is added (-value) For noble gases, energy is required to add an electron Electron Affinity Trends Little change down the group however, the electron affinity becomes more negative (decreases) across the period. 1 How is ionization energy compared to electron affinity? Ionic Radius The ion sizes of both the cation and anion in an ionic bond are used to calculate the bond energies between the atoms. The size of an ion depends on the ratio of nuclear charge to electrons. Cations are smaller than their neutral atoms Z > e - Anions are larger than their neutral atoms Z < e -

4 Observe an Isoelectronic Series Increasing Z eff O 2- F - Ne Na + Mg 2+ Decreasing size 1.Arrange the following in order of decreasing size: Se 2-, Sr 2+, Br -, Rb + Se 2- > Br - > [Kr] > Rb + > Sr 2+ Electronegativity Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself. Pauling Electronegativity Scale F = 4. Linus Pauling ( ) Electronegativity values, which are unit-less, are arbitrarily assigned beginning at 4. for F Electronegativity Trends Increases across the period and decreases down the group Electronegative atoms have a very negative electron affinity and a very large ionization energy Formation of Oxides Oxides of strong metallic elements are basic Oxides of strong non-metallic elements are acidic Oxides of other elements may be either acidic or basic (amphoteric) depending on environment

5 Occurrences in Nature Strong metals only exist in ionic compounds Weak metals exist as free elements and ionic compounds Non-metals exist in molecular elements, molecules and ionic compounds (exception C) Nobel gases exist as free atoms