Periodic trends valence electrons

Ionization lithium, 267 magnesium, 270 sodium, 270 Ionization energy, 267 alkaline earths, 379 and atomic number, 268 and ihe periodic table, 267 and valence electrons, 269 halogens, 353 measurement of, 268 successive, 269 table of, 268 trends, 268... 

Atomic radii typically decrease from left to right across a period and increase down a group (Fig. 14.2 see also Fig. 1.46). As the nuclear charge experienced by the valence electrons increases across a period, the electrons are pulled closer to the nucleus, so decreasing the atomic radius. Down a group the valence electrons are farther and farther from the nucleus, which increases the atomic radius. Ionic radii follow similar periodic trends (see Fig. 1.48). 

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. 

Polarizability shows periodic variations that correlate with periodic trends in how tightly valence electrons are bound to the nucleus ... 

On the basis of the Periodic Table, topics of intermetallic systematics will be presented in the next chapter. In the present chapter the Periodic Table will be revisited and its structure and subdivisions summarized. In relation also to some concepts previously presented, such as electronegativity, Mendeleev number, etc. described in Chapter 2, typical property trends along the Table will be shown. Strictly related concepts, such as Periodic Table group number, average group number and valence-electron number will be considered and used in the description and classification of intermetallic phase families. 

Comments on some trends and on the Divides in the Periodic Table. It is clear that, on the basis also of the atomic structure of the different elements, the subdivision of the Periodic Table in blocks and the consideration of its groups and periods are fundamental reference tools in the description and classification of the properties and behaviour of the elements and in the definition of typical trends in such characteristics. Well-known chemical examples are the valence-electron numbers, the oxidation states, the general reactivity, etc. As far as the intermetallic reactivity is concerned, these aspects will be examined in detail in the various paragraphs of Chapter 5 where, for the different groups of metals, the alloying behaviour, its trend and periodicity will be discussed. A few more particular trends and classification criteria, which are especially relevant in specific positions of the Periodic Table, will be summarized here. 

Ionization energy generally increases across a period. Again, this trend is linked to the atomic radius. Across a period, the atomic radius decreases because Zeff increases. The force of attraction between the nucleus and valence electrons is subsequently increased. Therefore, more energy is needed to remove one such electron. 

In addition to the information below, the periodic table shows trends in atomic size, ionization energy, electronegativity, valence electrons, and melting points. 

Metallic and nonmetallic properties are related to the number of valence electrons and the radius of an atom. Within a period, as the metallic properties decrease from left to right, the nonmetallic properties increase. Within a group as the metallic properties increase, the nonmetallic properties decrease from top to bottom. If the above trends are considered, francium, Fr, would be expected to have most metallic properties. However, since Fr is a radioactive element, not all of its properties have been determined yet. 

The outermost electrons, often called the valence electrons, are primarily responsible for the chemical properties of the elements. It follows that the elements in a specific group will show similar characteristic oxidation numbers (charges, also called valences) and display a trend in characteristics. Even though electron configurations were not known when the earliest periodic tables were formulated, the elements were placed by similarity of characteristics. 

Trends in ionization energy and electron affinity within a period reflect the stability of valence electron configurations. A stable system requires more energy to change and releases less when changed. Note the peaks in stability for groups 2, 13, and 16. 

As expected, the elements of a given column in the periodic table all behave in the same way because they all have the same number of valence electrons. There are a few exceptions to this rule, as you can tell if you look at the stairstep line passing through the periodic table. The elements to the right of the stairstep behave as nonmetals, whereas the elements to the left of this line behave as metals. But overall, the generalizations and trends apparent from the periodic table make it invaluable as a predictive tool. 

The ionization radii calculated for all atoms with this procedure show remarkable periodicity that mirrors many trends observed or inferred empirically for atomic properties such as electronegativity, covalent radii etc. Also indicated is a simple explanation for promotion of atoms into their valence state, before a chemical reaction commences [114]. This generally accepted mechanism, never satisfactorily explained before, can be accounted for simply in terms of environmental pressure. Whenever an atom is crowded because of high pressure or temperature or even concentration on a catalytic surface, the valence electron becomes promoted towards its ionization limit. In this limit the atom enters the valence state as an electron becomes decoupled from... 

Bismuth (Z = 83) is the heaviest stable element in group 15 (VA) of the periodic table (see Periodic Table Trends in the Properties of the Elements). The Bi isotope, which is 100% abundant, has a 9/2 nuclear spin. Bi, an alpha emitter is used in nuclear medicine as a radiotherapeutic agent. Bismuth has two stable oxidation states Bi(V), corresponding to complete loss of the valence electrons, and Bi(III), a lower oxidation state that retains two valence electrons. Both oxidation states are diamagnetic. The latter is more stable and more common since Bi(V) has a large reduction potential ... 

Figure 5-8 shows the variation of bond distance with the number of valence electrons in second-period p block homonuclear diatomic molecules. As the number of electrons increases, the number in bonding orbitals also increases, the bond strength becomes greater, and the bond length becomes shorter. This continues up to 10 valence electrons in N2 and then the trend reverses because the additional electrons occupy antibonding orbitals. The ions N2", 02, 02, and 0 are also shown in the figure and follow a similar trend. 

Trends within periods As shown in Figure 6-16 and by the values in Table 6-2, first ionization energies generally increase as you move left-to-right across a period. The increased nuclear charge of each successive element produces an increased hold on the valence electrons. 

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