Periodic property electron configurations

In so far as the chemical (and physical) properties of an element derive from its electronic configuration, and especially the configuration of its least tightly bound electrons, it follows that chemical periodicity and the form of the periodic table can be elegantly interpreted in terms of electronic structure. 

The alkali metals form a homogeneous group of extremely reactive elements which illustrate well the similarities and trends to be expected from the periodic classification, as discussed in Chapter 2. Their physical and chemical properties are readily interpreted in terms of their simple electronic configuration, ns, and for this reason they have been extensively studied by the full range of experimental and theoretical techniques. Compounds of sodium and potassium have been known from ancient times and both elements are essential for animal life. They are also major items of trade, commerce and chemical industry. Lithium was first recognized as a separate element at the beginning of the nineteenth eentury but did not assume major industrial importance until about 40 y ago. Rubidium and caesium are of considerable academic interest but so far have few industrial applications. Francium, the elusive element 87, has only fleeting existence in nature due to its very short radioactive half-life, and this delayed its discovery until 1939. 

The stability of the electronic configuration is indicated by the fact that each element has the highest ionization energy in its period, though the value decreases down the group as a result of increasing size of the atoms. For the heavier elements is it actually smaller than for first-row elements such as O and F with consequences for the chemical reactivities of the noble gases which will be considered in the next section. Nuclear properties, particularly for xenon, have been exploited for nmr spectroscopy and Mdssbauer... 

There are two immediate questions we ask about the transition elements once we know where they are in the periodic table (1) Why do we consider these elements together (2) What is special about their properties These questions are closely related because they both depend upon the electron configurations of the atoms. What, then, is the electron configuration we might expect for these elements ... 

The reason usually cited for the great similarity in the properties of the lanthanides is that they have similar electronic configurations in the outermost 6s and 5d orbitals. This occurs because, at this point in the periodic table, the added electrons begin to enter 4f orbitals which are fairly deep inside the atom. These orbitals are screened quite well from the outside by outer electrons, so changing the number of 4/electrons has almost no effect on the chemical properties of the atom. The added electrons do not become valence electrons in a chemical sense—neither are they readily shared nor are they readily removed. 

A primary goal of the periodic table is to assist recognition of the ground-state valence electron configuration of each atom, the chief determinant of its chemical properties" ([21], p 5). 

While it may be true that the periodic table is used in chemical education in order to arrive at the electronic configuration of any particular atom, this is surely not a primary goal of the periodic table for chemists in general. The main goal of the periodic table remains as the classification scheme for the properties of the elements, especially as they occur in chemical compounds.9... 

As in the discussion of hydrogen, in this section we examine the properties of the alkali metals in the context of the periodic table and focus on significant applications of the elements and selected compounds. The valence electron configuration of the alkali metals is s1, where n is the period number. Their physical and chemical properties are dominated by the ease with which the single valence electron can be removed (Table 14.3). 

The physical and chemical properties of the elements show regular periodic trends that can be explained using electron configurations and nuclear charges. We focus on the physical properties of the elements in this section. A preliminary discussion of the chemical properties of some of the elements appears in Section Other chemical properties are discussed after we introduce the principles of chemical bonding in Chapters 9 and 10. 

Ion formation is only one pattern of chemical behavior. Many other chemical trends can be traced ultimately to valence electron configurations, but we need the description of chemical bonding that appears in Chapters 9 and 10 to explain such periodic properties. Nevertheless, we can relate important patterns in chemical behavior to the ability of some elements to form ions. One example is the subdivision of the periodic table into metals, nonmetals, and metalloids, first introduced in Chapter 1. 

The periodic table orders the elements in a way that helps chemists understand why atoms behave as they do. What makes fluorine react violently with cesium while its nearest neighbor neon is reluctant to react with anything In other words, what gives the elements their properties and what order lies below the surface of their seemingly random nature Scientists know now that the periodicity of the elements is due largely to recurring patterns in their electron configurations. 

The previous chapter explored the elements—their electron configurations, their periodicity, and their properties. This chapter will investigate how chemists create more complex substances—the bits of matter called molecules. 

Periodic table A chart that arranges the elements by atomic number in a way that the vertical columns produce groups of elements with similar valence electron configurations and chemical properties. 

The goal of this chapter is to help you relate the properties of elements to their position on the periodic table. These properties include ionization energies and electron affinities. You may want to review the basic structure of the periodic table in Chapter 2 and electron configurations in Chapter 7. And Don t forget — Practice, Practice, Practice. 

Elements in a group have similar chemical properties because they have similar outer electron configurations. That is, they have the same number of valence electrons. This observation gives rise to three patterns that you can deduce from the periodic table. 

Electron Configurations, Atomic Properties, and Periodic Trends... 

In this section, you have seen how a theoretical idea, the quantum mechanical model of the atom, explains the experimentally determined structure of the periodic table, and the properties of its elements. Your understanding of the four quantum numbers enabled you to write electron configurations and draw orbital diagrams for atoms of the elements. You also learned how to read the periodic table to deduce the electron configuration of any element. 

The rare earth (RE) ions most commonly used for applications as phosphors, lasers, and amplifiers are the so-called lanthanide ions. Lanthanide ions are formed by ionization of a nnmber of atoms located in periodic table after lanthanum from the cerium atom (atomic number 58), which has an onter electronic configuration 5s 5p 5d 4f 6s, to the ytterbium atom (atomic number 70), with an outer electronic configuration 5s 5p 4f " 6s. These atoms are nsnally incorporated in crystals as divalent or trivalent cations. In trivalent ions 5d, 6s, and some 4f electrons are removed and so (RE) + ions deal with transitions between electronic energy sublevels of the 4f" electroiuc configuration. Divalent lanthanide ions contain one more f electron (for instance, the Eu + ion has the same electronic configuration as the Gd + ion, the next element in the periodic table) but, at variance with trivalent ions, they tand use to show f d interconfigurational optical transitions. This aspect leads to quite different spectroscopic properties between divalent and trivalent ions, and so we will discuss them separately. 

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