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«F_Name» «L_Name» Period «Per» Chapter 5 Section 3: Electron Configuration and Periodic Properties Book states You Say: Valence Electrons – The electrons available to be lost, gained, or shared in the formation of chemical compounds. Chemical compounds form because electrons are lost, gained, or shared between atoms. Main-Group Elements Main-group elements are the s-block and p-block elements (also known and the representative elements). Atomic Size (Atomic radius): is one-half the distance between the nuclei of identical atoms that are bonded together. Period Trend: atoms get smaller across a period This is caused by the increasing positive charge of the nucleus. Group Trend: In general, the atomic radii (size) of the MainGroup Elements increase down a group. The cause is that electrons occupy orbitals in higher energy levels Ion – an atom or group of bonded atoms that has a positive or negative charge. + Example: Sodium forms a Na ion. (cation) Ionization – Any process that result in the formation of an ion Ionization energy Ionization energy, IE (or first ionization energy, IE1) – is energy required to remove one electron from a neutral atom of an element. Ionization energy is measured in kilojoules per mole (kJ/mol) Period Trend: In general, ionization energies of the maingroup elements increase across each period. This increase is caused by increasing nuclear charge. Group Trend: among the main-group elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus. Plus there are electrons between the valence electrons and the nucleus that shield the outer electrons. The shielding makes valence electrons easier to remove. «Num» «F_Name» «L_Name» Period «Per» 2nd IE (IE2), or 3rd (IE3) Each successive electron removed from an ion has an increasingly stronger effective nuclear charge; therefore, they require more energy to remove. Electron Affinity – The energy change that occurs when an electron is added to a neutral atom; Energy released is represented by a negative number. Some atoms must be “forced” to gain an electron by the addition of energy. When energy is absorbed it is represented by a positive number and the ion is unstable and usually decays (breaks apart back to the element). Periodic Trend for Electron Affinity – It is easier to add electrons to the p block. An exception occurs between Groups 14 and 15. The ease with which halogen atoms gain electrons is a major reason for the high reactivities of the Group 17 elements. Group Trends for Electron Affinity – Trends for electron affinities within groups can be unpredictable; as a general rule, electrons are more difficult to add down a group. This pattern is a result of two competing factors: nuclear charge and size. In general, the size effect predominates and lowers the electron affinity. Ions A cation is a positive ion. An anion is a negative ion. Ionic Radii A cation forms by the loss of one or more electrons; This leads to a decrease in atomic radius because the removal of the highest-energy-level electrons causes a smaller electron cloud. An anion forms by the addition of one or more electrons; which always causes an increase in electron cloud size. Period Trends Within each period of the periodic table, the metals at the left tend to form cations and the nonmetals at the upper right tend to form anions. Group Trends Just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii. «Num» «F_Name» «L_Name» Period «Per» «Num» Sublevel Shielding – the inner electrons (d and f orbitals) shield the valence electrons causing the nucleus to have less pull on the valence electrons (Weakens the atoms hold on the valence electrons.) Electronegativity Electronegativity – measure of the ability of an atom in a chemical compound to attract electrons. Linus Pauling devised a scale of numerical values reflecting the tendency of an atom to attract electrons. The most electronegative element, fluorine, is arbitrarily assigned an electronegativity value of four. The rest of the element’s values are then calculated relative to fluorine Periodic Trend: Electronegativities increase across each period, although there are exceptions. Group Trend: Electronegativities either decrease down a group or remain about the same. The combination of the period and group trends in electronegativity results in the highest values belonging to the elements in the upper right of the periodic table. The lowest values belong to the elements in the lower left of the table. Noble Gases do not have assigned values – they don’t react readily. Periodic Properties of the d- and f-Block Elements The properties of the d-block elements vary less and with less regularity than those of the main-group elements. The d- and f-Block Elements Atomic Radii - The atomic radii of the d-block elements generally decrease across the periods, this decrease is less than that for the main-group elements because the electrons added to the (n − 1)d sublevel shield the valence electrons Atomic Radii Ionization Energy - ionization energies of the d-block and f-block elements generally increase across the periods. The first ionization energies of the dblock elements generally increase down each group. Ionization Energy «F_Name» «L_Name» Period «Per» «Num» Ion Formation and Ionic Radii - The order in which electrons are Ion Formation removed from all atoms of the d-block and f-block elements is exactly the and Ionic Radii reverse of the order given by the electron-configuration notation. In other words, electrons in the highest occupied sublevel are always removed first. For the d-block elements, this means that although newly added electrons that occupy the d sublevels are not removed first; the first electrons to be removed are those in the outermost s sublevels. For example, iron, Fe, has the electron configuration [Ar]3d64s2. It loses a 4s electron first to form Fe+ ([Ar] 3d 64s 1). Fe+ can then lose the second 4s electron to form Fe2+ ([Ar] 3d 6). Fe2+ can then lose a 3d electron to form Fe3+ ([Ar] 3d 5). Most d-block elements commonly form 2+ ions in compounds. Some, such as iron and chromium, also commonly form 3+ ions. The Group 3 elements form only ions with a 3+ charge. Copper forms 1+ and 2+ ions, and silver usually forms only 1+ ions. As expected, the cations have smaller radii than the atoms do. Comparing 2+ ions across the periods follows the decrease in size that parallels the decrease in atomic radii. Electronegativity - The d-block elements all have electronegativities between 1.1 and 2.54. Only the active metals of Groups 1 and 2 have lower electronegativities. The d-block elements also follow the general trend for electronegativity values to increase as radii decrease, and vice versa. The f-block elements all have similar electronegativities, which range from 1.1 to 1.5 Electronegativity