Earth Materials. John O'Brien

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СКАЧАТЬ a stable configuration and become divalent anions such as S−2. However, in the presence of highly electronegative oxygen these elements may lose electrons and become cations such as S+6.

       Column 17 (VIIA) nonmetallic elements such as Cl− and F−1 commonly exist as monovalent (−1) anions. Because electrons are very difficult to remove from their electron clouds, these elements tend to attract one additional electron into their highest principal quantum level to achieve a stable electron configuration.

       Column 18 (VIIIA) noble gas elements such as He, Ar, and Ne contain complete outer electron shells (s2, p6) and do not commonly combine with other elements to form minerals. Instead, they tend to exist as monatomic (composed of single atoms) gases.

      The periodic table is a highly visual and logical way in which to illustrate patterns in the electron configurations of the elements. Elements are grouped in rows or classes according to the highest principal quantum level in which electrons occur in the ground state. Elements are grouped into columns or groups based on similarities in the electron configurations in the higher principal quantum levels; those that are farthest from the nucleus and involved in most chemical reactions. More thorough explanations of the periodic table and the properties of elements are available in many standard texts in chemistry and physics.

      From the discussion above, it should be clear that during the chemical reactions that produce Earth materials, elements display behaviors that are related to their electron configurations. Group 18 (VIIIA) elements in the far right column of the periodic table have stable electron configurations and tend to exist as uncharged atoms. Metallic elements toward the left side of the periodic table are strongly electropositive and tend to give up one or more electrons to become positively charged particles called cations. Nonmetallic elements toward the right side of the periodic table, especially in groups 16 (VIA) and 17 (VIIA), are strongly electronegative and tend to attract electrons to become negatively charged particles called anions. Elements toward the middle of the periodic table are somewhat electropositive and tend to lose various numbers of electrons to become cations with various amounts of positive charge. These tendencies are summarized in Table 2.4.

Figure 2.6 Trends in variation of atomic radii (in angstroms; 1 Å = 10⁻¹⁰ m) with their position on the periodic table, illustrated by rows 3 and 4. With few exceptions, radii tend to decrease from left to right and from bottom to top.

      2.2.4 Atomic and ionic radii

Atomic radii are defined as half the distance between the nuclei of identical, bonded neighboring atoms. Because the electrons in higher quantum levels are farther from the nucleus, the effective radius of electrically neutral atoms generally increases from the top to bottom (row 1–row 7) in the periodic table (see Table 2.3). However, atomic radii generally decrease within rows from left to right (Figure 2.6). This occurs because the addition of electrons to a given quantum level does not significantly increase atomic radius, while the increase in the number of positively charged protons in the nucleus causes the electron cloud to contract as electrons are pulled closer to the nucleus. Atoms with large atomic numbers and large electron clouds include cesium (Cs), rubidium (Rb), potassium (K), barium (Ba), and uranium (U). Atoms with small atomic numbers and small electron clouds include hydrogen (H), beryllium (Be), and carbon (C).

Figure 2.7 Radii (in angstroms) of some common cations in relationship to the atomic radius of the neutral atoms.

      Electrons in the outer, higher energy electron levels are least tightly bound to the positively charged nucleus. This weak attraction results because these electrons are farthest from the nucleus and because they are shielded from the nucleus by the intervening electrons that occupy lower quantum level positions closer to the nucleus. These outer electrons or valence electrons are the electrons that are involved in a wide variety of chemical reactions, including those that produce minerals, rocks, and a wide variety of synthetic materials. The loss or gain of these valence electrons produces anions and cations, respectively.

When atoms are ionized by the loss or gain of electrons, their ionic radii, invariably change. This results from the electrical forces that act between the positively charged protons in the nucleus and the negatively charged electrons in the electron clouds. The ionic radii of cations tend to be smaller than the atomic radii of the same element (Figure 2.7). As electrons are lost from the electron cloud during cation formation, the positively charged protons in the nucleus tend to exert a greater force on each of the remaining electrons. This force draws electrons closer to the nucleus, reducing the effective radius of the electron cloud as it contracts. The larger the charge on the cation, the more its radius is reduced by the excess positive charge in the nucleus. This is well illustrated by the radii of the common cations of iron (Figure 2.8). Ferric iron (Fe⁺³) has a smaller radius (0.64 Å) than does ferrous iron (Fe⁺² = 0.74 Å). Both iron cations possess much smaller radii than neutral iron (Fe⁰ = 1.23 Å) in which there is no excess positive charge in the nucleus.

Figure 2.8 Radii (in angstroms) of some common anions in relationship to the atomic radius of the neutral atoms.

The ionic radii of anions are significantly larger than the atomic radii of the same neutral (uncharged) element (Figure 2.8). When electrons are added to the electron cloud during anion formation, the positively charged protons in the nucleus exert a smaller force on each of the electrons. This allows the electrons to move farther away from the nucleus, which causes the electron cloud to expand, increasing the effective radius of the anion. The larger the charge on the anion, the more its effective radius is increased.

Figure 2.9 Radii (in angstrom units) of some common anions and cations of sulfur in relationship to the neutral atom radius.

The expansion of anions and the contraction of cations are well illustrated by the common ions of sulfur (Figure 2.9). The divalent sulfur (S⁻²) anion possesses a relatively large average radius of 1.84 Å. In this case, the two electrons gained during the formation of a divalent sulfur anion produce a large deficit between positive charges in the nucleus and negative charges in the electron cloud. This leads to a significant increase in the effective ionic radius compared to that of electrically neutral sulfur (S⁰ СКАЧАТЬ